GIFT  OF 
Dean  Frank  H.  Probert 


Mining  Dept 


AMERICAN  SCIENCE  SERIES— ADVANCED   COURSE 


INORGANIC  CHEMISTRY 


BY 


IEA   REMSEN 

Professor  of  Chemistry  in  the  Johns  Hopkins  University 


FIFTH  EDITION,    REVISED 


NEW  YORK 

HENRY   HOLT   AND   COMPANY 

1899 


MM* 

OEPI. 


QlffI 

DEAN  FRANK  H  f 

DEPT. 


Copyright   1889,  1898, 

LLi       V/t:/: 

HENRY  HOLT'«evC 


ROBERT  DRUMMOND,  Printed  by 

Electrotyper,  D   G.  F,  CLASS, 

NEW  YORK.  NEW  YORK. 


PREFACE  TO  THE  FIRST  EDITION. 


IN  the  preparation  of  this  book  I  have  been  much 
encouraged  by  the  cordial  reception  which  has  been 
given  my  earlier  text-books,  both  in  this  country  and 
abroad.'  While  those  earlier  works  are  intended  to  form 
a  series  of  which  the  present  volume  is  the  most  advanced 
member,  it  has  little  in  common  with  the  lower  mem- 
bers except  the  general  method  of  treatment.  An  occa- 
sional paragraph  from  the  "Briefer  Course"  has  been 
incorporated,  but  the  two  books  are  quite  distinct. 

In  classifying  the  elements  the  periodic  system  has 
been  adopted,  and  this  has  been  pretty  closely  adhered 
to.  In  order  to  secure  as  logical  treatment  as  possible 
it  has  been  thought  best  not  to  give  detailed  descriptions 
of  apparatus  and  specific  directions  for  the  preparation 
of  substances,  in  the  text  proper.  By  avoiding  these  the 
attention  can  be  better  directed  to  the  principles  involved 
and  a  clearer  conception  of  these  principles  will  be  formed, 
than  when  the  attention  is  distracted  by  the  reading  of  such 
details.  On  the  other  hand,  full  descriptions  of  apparatus 
and  processes  will  be  found  in  the  Appendix  ;  and  these, 
it  is  believed,  will  be  of  service  to  the  teacher  in  the 
lecture-room,  as  well  as  to  the  student  in  the  laboratory. 

The  feature  of  the  book  which  perhaps  most  distin- 
guishes it  from  others  is  the  fulness  with  which  general 
relations  are  discussed  in  it.  Attention  is  constantly 
called  to  analogies  between  properties  of  substances  and 
between  chemical  reactions,  so  that  the  thoughtful 
student  will,  it  is  hoped,  be  led  to  look  upon  the  sub- 
stances and  the  reactions  not  as  independent  of  one 
another,  but  as  related  in  many  ways,  and  thus  forming 

iii 


ir  PREFACE. 

parts  of  a  system.  All  thinking  chemists  have  no  doubt  at 
times  an  indistinct  vision  of  a  perfect  science  of  Chemis- 
try yet  to  come,  in  which  the  relations  of  the  parts  will 
be  clearly  seen,  and  in  which  much  that  now  appears  of 
little  or  no  importance  will  be  recognized  as  significant. 
The  subject  cannot  as  yet,  however,  be  treated  as  if  that 
perfection  had  been  reached.  Much  progress  has  been 
made  of  late  years  in  the  classification  of  the  facts,  and 
it  is  of  prime  importance  to  the  student  that  general  rela- 
tions should  be  pointed  out  for  him  as  clearly  as  possible. 
Of  course  in  the  classification  of  facts  the  end  is  not 
reached.  In  every  case  of  chemical  action  there  are 
certain  features  which  call  for  much  deeper  study  than 
is  usually  given  to  them.  For  the  most  part  chemists 
have  been  content  to  know  what  chemical  changes  take 
place  when  two  or  more  substances  are  brought  into 
action,  and  have  paid  much  less  attention  to  the  accom- 
panying phenomena ;  and  yet  it  is  evident  that,  in 
order  to  get  a  clear  conception  of  the  nature  of  the 
chemical  act,  it  is  necessary  that  we  should  learn  all 
we  possibly  can  in  regard  to  that  act.  Of  late  years 
more  and  more  attention  has  been  given  to  the  study  of 
the  phenomena  accompanying  chemical  changes  ;  and 
a  clearer  view  has  been  gained  regarding  chemical 
action.  A  great  field  of  study  is  thus  opened,  which 
bears  to  the  science  of  Chemistry  as  a  whole  somewhat 
the  same  relation  that  Physiology  bears  to  Biology,  while 
the  study  of  chemical  substances  and  their  changes  as 
usually  carried  on  is  in  the  same  way  the  counterpart 
of  Morphology.  Neither  of  the  parts  taken  separately 
is  Chemistry  in  the  fullest  sense.  It  will  never  be  pos- 
sible to  study  Chemistry  without  taking  up  and  working 
with  chemical  substances ;  but  as  knowledge  grows, 
more  and  more  attention  will  surely  be  given  to  chemical 
action.  In  this  book  considerable  space  is  devoted 
to  the  discussion  of  the  results  obtained  in  the  latter 
kind  of  study.  Some,  no  doubt,  will  hold  that  even 
more  prominence  should  have  been  given  to  this  side  of 
the  subject.  Indeed  I  shall  be  glad  if  some  of  those  who 
use  the  book  become  interested  in  the  new  problems, 


PREFACE.  V 

and  go  further  into  their  study.  It  has  not,  however, 
appeared  to  me  advisable,  considering  the  purposes  for 
which  this  book  has  been  written,  to  discuss  them  more 
fully. 

The  subject  of  the  Constitution  of  Chemical  Compounds 
receives  a  due  share  of  attention.  Constitutional  for- 
mulas are  not,  however,  used  recklessly  as  though  they 
were  provided  by  nature  ready-made ;  but  the  effort  is 
made  to  keep  clearly  in  mind  the  facts  which  they  ex- 
press so  that  they  may  be  used  intelligently.  In  this 
connection  I  may  call  special  attention  to  the  way  in 
which  the  constitution  of  the  so-called  double  salts  of 
the  halogens  is  treated.  To  those  who  have  not  care- 
fully looked  into  the  evidence,  the  formulas  used  will 
perhaps  appear  too  speculative.  I  should  be  sorry  to 
err  in  this  direction.  For  some  time  past  the  view  put 
forward  has  seemed  to  me  to  be  justified,  and  I  find  that 
others  whose  judgment  I  respect  have  held  the  same 
view  at  least  in  regard  to  some  of  the  compounds  in 
question.  As,  generally  speaking,  these  compounds  are 
treated  inadequately,  and  as  they  are  commonly  regarded 
as  inexplicable,  I  propose  soon  to  present,  in  the  proper 
place,  the  evidence  upon  which  my  present  view  rests, 
when  it  will,  I  think,  be  found  that  the  evidence  is  fully 
as  strong  as  that  upon  which  our  views  concerning  the 
constitution  of  most  compounds  are  founded. 

IRA  EEMSEN, 
BALTIMORE,  March,  1889. 


PREFACE  TO  SECOND  EDITION. 


THE  call  for  a  new  edition  of  this  book  has  given  me 
an  opportunity  to  make  some  desirable  changes,  and  to 
correct  those  errors  to  which  my  attention  has  been 
directed  by  others  or  which  I  have  myself  discovered. 
The  revision  is  based  upon  the  labors  of  a  very  consid- 


VI  PREFACE. 

erable  number  of  readers  who  have  given  me  the  benefit 
of  their  criticisms,  and  I  take  this  opportunity  to  express 
my  sincere  thanks  to  all  those  who  have  aided  me. 
Should  any  one  using  the  new  edition  discover  errors  in 
it,  I  shall  be  thankful  to  be  informed  of  the  fact.  It 
seems  fair  to  say  that  I  have  heard  only  words  of  com- 
mendation in  regard  to  the  general  plan  and  spirit  of 
the  book. 

IRA  EEMSEN. 
BALTIMORE,  Decembw  6,  1889. 


PREFACE  TO  FIFTH  EDITION. 


DURING  the  eight  years  that  have  passed  since  the 
second  edition  of  this  book  was  published  a  number  of 
minor  corrections  have  been  made  in  it  from  time  to 
time.  It  has  now,  however,  been  subjected  to  a  thorough 
revision,  and  it  is  hoped  that  this  new  edition  will  be 
found  to  contain  everything  that  can  fairly  be  looked  for 
in  a  book  of  its  size.  A  new  Appendix  has  been  added 
containing  much  information  concerning  the  properties 
of  a  large  number  of  compounds  which  are  necessarily 
treated  briefly  or  not  at  all  in  the  text. 

It  may  not  be  inappropriate  to  mention  the  fact  that 
the  book  has  been  well  received  not  only  in  this  country, 
but  in  England  and,  in  Germany,  a  German  translation 
having  appeared  shortly  after  the  publication  of  the  first 
American  edition. 

IRA  EEMSEN. 

BALTIMORE,  February  21,  1898. 


CONTENTS. 


CHAPTER  I. 

CHEMICAL  AND  PHYSICAL  CHANGE — EARLIEST  CHEMICAL  KNOWLEDGE— 
LAW  OF  THE  INDESTRUCTIBILITY  OF  MATTER — LAW  OF  DEFINITE 
PROPORTIONS — LAW  OF  MULTIPLE  PROPORTIONS— THE  ELEMENTS. 

PAGE 

Matter  and  Energy — Chemical  Change — Physical  Change — Physics 
and  Chemistry — Earliest  Chemical  Knowledge — Alchemy — 
Chemistry  as  a  Science — Lavoisier's  Work — Law  of  the  Inde- 
structibility of  Matter  —  Conservation  of  Energy  —  Early 
Views  regarding  the  Composition  of  Matter — Elements — Chem- 
ical Action — Chemical  Affinity — Chemical  Compounds  and 
Mechanical  Mixtures — Qualitative  and  Quantitative  Study  of 
Chemical  Changes — Law  of  Definite  Proportions — Law  of  Mul- 
tiple Proportions — Combining  Weights  of  the  Elements — The 
Elements,  their  Symbols  and  Atomic  Weights— Symbols  of 
Compounds— Chemical  Equations — The  Scope  of  Chemistry — 
Chemical  Action  accompanied  by  other  Kinds  of  Action,  .  .  1 


CHAPTER  II. 

A  STUDY  OF  THE  ELEMENT  OXYGEN. 

Historical— Occurrence— Preparation— Physical  Properties— Chem- 
ical Properties — Burning  in  the  Air  and  Burning  in  Oxygen — 
Phlogiston  Theory — Lavoisier's  Explanation  of  Combustion — 
Kindling  Temperature— Slow  Oxidation— Heat  of  Combustion 
— Heat  of  Decomposition— Chemical  Energy  and  Chemical 
Work-Oxides, 28 

CHAPTER  III. 

A  STUDY  OF  THE  ELEMENT  HYDROGEN. 

Historical — Occurrence — Preparation— Physical  Properties— Chem- 
ical Properties — Comparison  of  Oxygen  and  Hydrogen,  ...  40 

vii 


viii  CONTENTS. 

CHAPTER  IV. 

STUDY  OP  THE  ACTION  OP  HYDROGEN  ON  OXYGEN. 

PAGB 

Burning  of  Hydrogen — Method  of  Dumas— Eudiometric  Method — 
Calculation  of  the  Results  obtained  in  exploding  Mixtures  of 
Hydrogen  and  Oxygen — Determination  of  tho  Volume  of  Water 
Vapor  formed  by  Union  of  Definite  Volumes  of  Hydrogen  and 
Oxygen — Heat  evolved  in  the  Union  of  Hydrogen  and  Oxy- 
gen— Applications  of  the  Heat  formed  by  the  Combination  of 
Hydrogen  and  Oxygen — Oxyhydrogen  Light— Velocity  of 
Combination  of  a  Mixture  of  Hydrogen  and  Oxygen — Sum- 
mary,   49 

CHAPTER  V. 

WATER. 

Historical — Occurrence — Formation  of  Water  and  Proofs  of  its 
Composition — Properties  of  Water — Chemical  Properties  of 
Water — Water  as  a  Solvent — Solution  as  an  Aid  to  Chemical 
Action — Natural  Waters— What  constitutes  a  Bad  Drinking 
Water— Purification  of  Water, 57 

CHAPTER  VI. 

CONSTITUTION  OP  MATTER — ATOMIC  THEORY — ATOMS  AND  MOLECULES- 
CONSTITUTION— VALENCE. 

Early  Views— The  Atomic  Theory  as  proposed  by  Dalton — Use  and 
Value  of  a  Theory — Atomic  Weights  and  Combining  Weights 
— Molecules — Avogadro's  Law — Distinction  between  Molecules 
and  Atoms  —  Molecular  Weights  —  Deduction  of  Atomic 
Weights  from  Molecular  Weights — Exact  Atomic  Weights 
determined  by  the  Aid  of  Analysis — Molecular  Formulas — 
Constitution— Valence— Replacing  Power  of  Elements,  ...  68 


CHAPTER  VII. 

OZONE — ALLOTROPY — NASCENT   STATE — HYDROGEN   DIOXIDE. 

Occurrence— Preparation— Properties— Relation  between  Oxygen 
and  Ozone — Ozone  in  the  Air — Allotropy — Varying  Number 
of  Atoms  in  the  Molecules  of  one  and  the  same  Element — 
Nascent  State —Hydrogen  Dioxide  or  Hydrogen  Peroxide — 
Properties — Occurrence  in  the  Air— Characteristic  Reactions — 
Thermochemical  Considerations,  .  ...  85 


CONTENTS.  ix 

CHAPTER  VIII. 

CHLORINE — HYDROCHLORIC   ACID. 

PAGE 

Historical— Occurrence  of  Chlorine— Preparation — Weldon's  Proc- 
ess—Electrolytic Process — Properties — Different  Kinds  of 
Action — Chlorine  Hydrate  and  Liquid  Chorine — Applications 
of  Chlorine — Hydrochloric  Acid — Historical — Study  of  the 
Action  of  Hydrogen  upon  Chlorine— Preparation — Properties 
—Chemical  Action  of  Hydrochloric  Acid, 96 

CHAPTER  IX. 

COMPOUNDS  OP  CHLORINE  WITH  OXYGEN  AND  WITH  HYDROGEN  AND 

OXYGEN. 

General — Principal  Reactions  for  Making  Compounds  of  Chlorine 
with  Hydrogen  and  Oxygen — Chloric  Acid — Properties— 
Hypochlorous  Acid — Chlorous  Acid — Perchloric  Acid — Gen- 
eral— Compounds  of  Chlorine  with  Oxygen — Constitution  of 
the  Compounds  of  Chlorine  with  Hydrogen  and  Oxygen — 
Comparison  of  Chlorine  and  Oxygen, 113 

CHAPTER  X. 

ACIDS — BASES — NEUTRALIZATION — SALTS. 

General — A  Study  of  the  Act  of  Neutralization— General  Statements 
— Definitions — Comparison  of  the  Reaction  between  Acids  and 
Hydroxides,  and  between  Acids  and  Chlorides — Other  Similar 
Reactions— Distinction  between  Acids  and  Bases— Metals  or 
Base-forming  Elements — Constitution  of  Acids  and  Bases — 
Constitution  of  Salts— Basicity  of  Acids— Acidity  of  Bases- 
Salts — Acid  Properties  and  Oxygen — Nomenclature  of  Acids — 
Nomenclature  of  Bases — Nomenclature  of  Salts, 127 

CHAPTER  XI. 

NATURAL  CLASSIFICATION  OF  THE  ELEMENTS— THE  PERIODIC  LAW. 

Historical— Arrangement  of  the  Elements— Connection  between 
the  Position  of  the  Elements  in  the  Natural  System  and  their 
Chemical  Properties— Plan  to  be  followed, 147 

CHAPTER  XII. 

THE   ELEMENTS   OF   FAMILY  VII,    GROUP   B: 
FLUORINE— CHLORINE — BROMINE— IODINE. 

General— Bromine — Occurrence — Preparation— Properties — Chem- 
ical Conduct  of  Bromine — Uses  of  Bromine — Hydrobromic 


CONTENTS. 

PACK 

Acids — Properties — Compounds  of  Bromine  with  Hydrogen  and 
Oxygen — Compounds  of  Bromine  and  Chlorine — Iodine — 
Occurrence — Preparation — Properties — Hydriodic  Acid — lodic 
Acid — Iodine  Pentoxide  or  lodic  Anhydride— Anhydrides,  or 
Acidic  Oxides— Periodic  Acid — Periodates— Constitution  of 
Periodic  Acid — Constitution  of  lodic  Acid  and  the  Oxygen 
Acids  of  Bromine — Compounds  of  Iodine  with  Chlorine — 
Compounds  of  Iodine  with  Bromine — Fluorine — Occurrence — 
Properties— Hydrofluoric  Acid— Constitution  of  Hydrofluoric 
Acid  and  the  Fluorides — Compound  of  Fluorine  with  Iodine- 
Tabular  Presentation  of  the  Compounds  of  the  Members  of  the 
Chlorine  Family  with  Hydrogen,  with  Oxygen,  with  Hydrogen 
and  Oxygen,  and  with  One  Another — Relative  Affinities  of  the 
Elements  of  the  Chlorine  Group — Family  VII,  Group  A — 

.  160 


CHAPTER  XIII. 

THE   ELEMENTS   OP   FAMILY  VI,    GROUP   B  : 
SULPHUR — SELENIUM — TELLURIUM. 

Introductory — Sulphur— Occurrence— Extraction  of  Sulphur  from 
its  Ores — Properties — Uses  of  Sulphur — Compounds  of  Sulphur 
with  Hydrogen — Hydrogen  Sulphide,  Sulphuretted  Hydrogen 
— Properties — Action  of  Hydrogen  Sulphide  upon  Solutions  of 
Salts,  Use  in  Chemical  Analysis — Hydrosulphides — Hydrogen 
Persulphide — Compounds  of  Sulphur  with  Members  of  the 
Chlorine  Group — Selenium — Occurrence — Properties— Hydro- 
gen Selenide  — Tellurium  —  Occurrence — Properties — Hydro- 
gen Telluride, 185 

CHAPTER    XIV. 

COMPOUNDS  OF  SULPHUR,  SELENIUM,  AND  TELLURIUM  WITH  OXYGEN 
AND  WITH  OXYGEN  AND  HYDROGEN. 

Introductory — Sulphuric  Acid — Pure  Sulphuric  Acid— Tetrahy- 
droxyl  Sulphuric  Acid — Normal  Sulphuric  Acid — Disulphuric 
Acid,  Pyrosulphuric  Acid — Sulphurous  Acid — Hyposulphurous 
Acid— Thiosulphuric  Acid— Other  Acids  of  Sulphur— Per- 
sulphuric  Acid — Constitution  of  the  Acids  of  Sulphur— Com- 
pound of  Sulphur  with  Oxygen — Sulphur  Dioxide — Sulphur 
Trioxide — Acid  Chlorides  of  Sulphur — Thionyl  Chloride — 
Sulphtiryl  Chloride  —  Chlorsulphuric  Acid,  or  Sulphuryl- 
hydroxyl  Chloride— Compounds  of  Selenium  and  Tellurium 
with  Oxygen  and  with  Oxygen  and  Hydrogen — Selenious  Acid 
— Selenic  Acid — Selenium  Dioxide — Acid  Chlorides  of  Sele- 
nium—Tellurious  Acid— Telluric  Acid— Oxides  of  Tellurium— 
Sulphotelluric  Acid— Family  VI,  Group  A .  .206 


CONTENTS.  xi 

CHAPTER  XV. 

NITEOGEN — THE  AIK— ARGON. 

PAGE 

Nitrogen — General— Occurrence  of  Nitrogen — Preparation — Prop- 
erties— The  Air — Analysis  of  Air — Argon, 248 

CHAPTER  XVI. 

COMPOUNDS    OF    NITROGEN    WITH    HYDROGEN — WITH    HYDROGEN    AND 
OXYGEN— WITH   OXYGEN,    ETC. 

General  Conditions  which  give  Rise  to  the  Formation  of  the  Sim- 
pler Compounds  of  Nitrogen — RelatioDs  between  the  Principal 
Compounds  of  Nitrogen — Ammonia — Composition  of  Am- 
monia— Ammonium  Amalgam — Metallic  Derivatives  of  Am- 
monium Compounds  and  of  Ammonia — Structure  of  Ammoni- 
um Compounds — Hydraziue — Hydroxylamine — Triazoic  Acid 
— Nitric  Acid — Red  Fuming  Nitric  Acid — Nitrous  Acid — Hy- 
ponitrous  Acid— Nitrous  Oxide— Nitric  Oxide— Nitrogen  Tri- 
oxide — Nitrogen  Peroxide — Nitrogen  Pentoxide — Structure  of 
the  Compounds  of  Nitrogen  with  Oxygen  and  Hydrogen — 
Compounds  of  Nitrogen  with  the  Elements  of  the  Chlorine 
Group — Compounds  of  Nitrogen  with  the  Members  of  the 
Sulphur  Group, 260 

CHAPTER  XVII. 

ELEMENTS   OP  FAMILY  Y,    GROUP   B: 

PHOSPHORUS — ARSENIC— ANTIMONY— BISMUTH.        THE     ELEMENTS     AND 
THEIR   COMPOUNDS  WITH   HYDROGEN. 

General— Phosphorus— Occurrence — Preparation — Properties— Ap- 
plications of  Phosphorus — Compounds  of  Phosphorus  with 
Hydrogen — Phosphine,  Gaseous  Phosphuretted  Hydrogen- 
Arsenic — Occurrence — Preparation  —  Properties  — Arsine,  Ar- 
seniuretted  Hydrogen — Antimony — Occurrence — Properties — 
Applications  of  Antimony — Stibine — Methods  of  distinguishing 
between  Arsenic  and  Antimony — Bismuth— Occurrence— Com- 
pounds of  the  Members  of  the  Phosphorus  Group  with  the 
Members  of  the  Chlorine  Group — Phosphorus  Trichloride — 
Phosphorus  Pentachloride— Arsenic  Trichloride— Compounds 
of  Antimony  and  Chlorine — Bismuth  and  Chlorine — Double 
Salts, 294 

CHAPTER  XVIII. 

COMPOUNDS  OF  THE  ELEMENTS  OF  THE  PHOSPHORUS  GROUP  WITH 
OXYGEN  AND  WITH  OXYGEN  AND  HYDROGEN. 

Introduction — Phosphoric  Acid,  Orthophosphoric  Acid— Proper- 
ties— Pyrophosphoric  Acid — Metaphosphoric  Acid — Phosphor- 


xii  CONTENTS. 

PAGE 

ous  Acid — Hypophosphoric  Acid— Hypopliosphorous  Acid- 
Phosphorus  Peutoxide,  Phosphoric  Anhydride — Phosphorus 
Trioxide  or  Phosphorous  Anhydride — Phosphorus  Suboxide — 
Phosphorus  Tetroxide — Constitution  of  the  Acids  of  Phos- 
phorus— Phosphorus  Oxychloride — Arsenic  Acid— Arsenious 
Acid  —  Arsenic  Trioxide  —  Arsenic  Pentoxide  —  Sulphides  — 
Arsenic  Disulphide — Arsenic  Trisulphide — Arsenic  Pentasul- 
phide — Antimonic  Acid — Antimony  Trioxide— Salts  of  Anti- 
mony— Antimony  Tetroxide — Antimony  Peutoxide— Antimony 
Trisulphide  —  Antimony  Pentasulphide  —  Constitution  of  the 
Acids  of  Arsenic  and  Antimony — Oxychlorides  of  Antimony 
— Oxides  of  Bismuth — Salts  of  Bismuth — Bismuth  Dioxide — 
Bismuth  Peutoxide  —  Bismuth  Trisulphide  —  Bismuth  Oxy- 
chloride— Family  V,  Group  A— Vanadium — Vauadic  Acid — 
Tantalum  —  Columbium  —  Didymium — Boron — General  — Oc- 
currence— Preparation — Properties — Boron  Trichloride — Boron 
Trifluoride — Boric  Acid — Salts  of  Boron — Nitrogen  Boride,  .  321 

CHAPTER  XIX. 

CAKBON    AND    ITS    SIMPLER    COMPOUNDS    WITH  HYDROGEN  AND  CHLO- 
RINE. 

Introductory — Occurrence  of  Carbon — Diamond— Graphite — Amor- 
phous Carbon — Coal — Diamond,  Graphite,  and  Charcoal  are 
Different  Forms  of  the  Element  Carbon — Chemical  Conduct  of 
Carbon — Compounds  of  Carbon  with  Hydrogen,  or  Hydrocar- 
bons. Conditions  under  which  Hydrocarbons  are  formed — 
Number  of  Hydrocarbons— Homology,  Homologous  Series- 
Cause  of  the  Homology  among  Compounds  of  Carbon — Other 
Series  of  Hydrocarbons — Marsh  Gas,  Methane,  Fire-damp — 
Ethylene,  Olefiant  Gas— Acetylene — Simpler  Compounds  of 
Carbon  with  the  Members  of  the  Chlorine  Group, 357 

CHAPTER  XX. 

SIMPLER  COMPOUNDS  OF  CARBON  WITH  OXYGEN,  AND  WITH  OXYGEN 
AND  HYDROGEN. 

General — Relations  between  the  Compounds  of  Carbon  with  Hy- 
drogen and  Oxygen — Carbon  Dioxide — Preparation — Proper- 
ties—Relations of  Carbon  Dioxide  to  Chemical  Energy — 
Respiration — Carbon  Dioxide  and  Life — Energy  Stored  up  in 
Plants — Carbonic  Acid  and  Carbonates — Carbon  Monoxide — 
Formic  Acid— Carbonyl  Chloride,  Phosgene, 376 

CHAPTER  XXI. 

ILLUMINATION — FLAME — BLOW-PIPE. 
COMPOUNDS   OF  CARBON   WITH  NITROGEN   AND   SULPHUR. 

Introduction  — Illuminating  Gas,  Coal  Gas  —  Flames  —  Kindling 
Temperature  of  Gases — Miner's  Safety-lamp— Structure  of 


CONTENTS.  xiii 

PAGE 

Flames— Blow-pipe— Causes  of  the  Luminosity  of  Flames— 
Bunsen  Burner— Compounds  of  Carbon  with  Nitrogen  and 
with  Sulphur— Cyanogen— Hydrocyanic  Acid,  Prussic  Acid- 
Cyanic  Acid— Carbon  Bisulphide— Sulphocarbonic  Acid,  Thio- 
carbonic  Acid— Oxysulphides— Sulphocyanic  Acid— Constitu- 
tion of  Cyanogen  and  its  Simpler  Compounds, 394 

CHAPTER  XXII. 

ELEMENTS   OP   FAMILY  IV,  GROUP   A  : 
SILICON— TITANIUM — ZIRCONIUM — CERIUM— THORIUM. 

General— Silicon — Occurrence  —  Preparation  —  Silicon  Hydride — 
Titanium— Zirconium — Thorium— Cerium— Compounds  of  the 
Elements  of  the  Silicon  Group  with  those  of  the  Chlorine 
Group— Silicon  Tetrachloride — Silicon  Hexachloride— Silicon 
Tetrafluoride — Constitution  of  Fluosilicic  Acid — Titanium  Tet- 
rachloride—Titanium  Tetrafluoride — Zirconium  Tetrachloride 
— Thorium  Tetrachloride — Thorium  Tetrafluoride — Compari- 
son of  the  Chlorides  of  Family  IV  with  those  of  Fainity  V — 
Compounds  of  the  Members  of  the  Silicon  Group  with  Oxygen, 
and  with  Oxygen  and  Hydrogen— Silicon  Dioxide— Properties 
—Uses— Silicic  Acid  —  Polysilicic  Acids  —  Disilicic  Acids— 
Trisilicic  Acids — Titanium  Dioxide — Zirconium  Dioxide — 
Thorium  Dioxide— Silicides— Family  IV,  Group  B,  .  .  .  .409 

CHAPTER  XXIII. 

CHEMICAL   ACTION. 

Retrospective— Classification  of  Reactions  of  the  Elements  and 
Compounds  Studied— Kinds  of  Chemical  Reactions— Direct 
Combination — Direct  Decomposition— Metathesis— The  Cause 
of  Chemical  Reactions — An  Ideal  Chemical  Reaction — Influ- 
ence of  Mass— Reactions  may  be  complete  if  one  of  the  Prod- 
ucts formed  is  Insoluble  or  Volatile— Thermochemical  Study 
of  Aflinity — Value  of  Thermochemical  Measurements — Heat 
of  Neutralization — Avidity  of  Acids — Other  Methods  for  De- 
termining the  Avidity  of  Acids— Study  of  Chemical  Decom- 
positions— Dissociation — Electrolysis — Electrolytic  Dissociation 
— Relations  between  Specific  Heat  and  Atomic  Weights — Ex- 
ceptions to  the  Law  of  Specific  Heats— Raoult's  Method  for 
the  Determination  of  Molecular  Weights — Determination  of 
the  Extent  of  Dissociation  of  a  Dissolved  Substance,  ....  426 

CHAPTER  XXIV. 

BASE-FORMING   ELEMENTS — GENERAL   CONSIDERATIONS. 

Introductory— Metallic  Properties — Order  in  which  the  Base-form- 
ing Elements  will  be  taken  up — Occurrence  of  the  Metals — 
Extraction  of  the  Metals  from  their  Ores— The  Properties  of 


Xiv  CONTENTS. 

PAGE! 

the  Metals — Compounds  of  the  Metals— Chlorides — Forma- 
tion of  Salts  in  General — General  Properties  of  the  Chlorides 
— The  so-called  Double  Chlorides  and  similar  Compounds  of 
Fluorine,  Bromine,  and  Iodine — Different  Chlorides  of  the 
same  Metal — Oxides — Differen-t  Oxides  of  the  same  Metal — 
Hydroxides — Decomposition  of  Salts  by  Bases— Sulphides — 
Hydrosulphides  —  Sulpho-salts  —  Nitrates  —  Chlorates  —  Sul- 
phates—Carbonates— Phosphates — Silicates, 455 

CHAPTER  XXV. 

ELEMENTS   OF  FAMILY   I,    GROUP   B  : 

THE      ALKALI    METALS  : — LITHIUM — SODIUM — POTASSIUM — RUBIDIUM — 
CvESIUM — AKMONIUM. 

General— Potassium — Occurrence — Preparation  —  Properties  — Po- 
tassium Hydride — Potassium  Fluoride,  Chloride,  Bromide, 
Iodide  —  Properties  — Applications  —  Potassium  Hydroxide — 
Potassium  Oxide — Potassium  Hydrosulpbide — Potassium  Sul- 
phide— Potassium  Nitrate — Applications — Gunpowder — Potas- 
sium Nitrite — Potassium  Chlorate — Potassium  Perchlorate — 
Potassium  Periodate — Potassium  Cyanide — Potassium  Cyanate 
— Potassium  Sulphocyanate — Potassium  Sulphate — Primary, 
or  Acid,  Potassium  Sulphate — Sulphites — Carbonates — Acid 
Potassium  Carbonate — Phosphates — Potassium  Silicate — Rubi- 
dium— Caesium — Sodium — Occurrence — Preparation —  Proper- 
ties— Applications — Sodium  Hydride— Sodium  Chloride — So- 
dium Hydroxide — Oxides — Sodium  Peroxide — Sodium  Sul- 
phantimonate — Sodium  Nitrate — Sodium  Sulphate — Sodium 
Thiosulphate — Sodium  Carbonate — Properties — Applications — 
The  Le  Blanc  Process  for  the  Manufacture  of  Sodium  Carbon- 
ate— Ammonia  Process  for  the  Manufacture  of  Soda — Manu- 
facture of  Soda  from  Cryolite — Mono-Sodium  Carbonate,  Pri- 
mary Sodium  Carbonate — Sodium-Potassium  Carbonate — 
Phosphates  —  Sodium  Metaphosphate  —  Di-sodium  Pyro-anti- 
monate — Sodium  Borate — Sodium  Silicate — Lithium — Lithium 
Phosphate — Lithium  Carbonate — Lithium  -Chloride — Ammo- 
nium Salts — Ammonium  Chloride — Ammonium  Sulphocyan- 
ate— Ammonium  Sulphide — Ammonium  Nitrate — Ammonium 
Carbonate — Sodium-ammonium  Phosphate— Reactions  of  the 
Members  of  the  Sodium  Group  which  are  of  Value  in  Chemical 
Analysis— Flame  Reactions  and  the  Spectroscope, 482 

CHAPTER   XXVI. 

ELEMENTS   OF   FAMILY  II,    GROUP   A  : 
GLUCINUM — MAGNESIUM — CALCIUM — STRONTIUM — BARIUM  [ERBIUM]. 

General — Calcium  Sub-Group  — Calcium — Occurrence  —  Prepara- 
tion— Properties— Calcium  Chloride— Calcium  Fluoride— Cal- 
cium Oxide— Calcium  Hydroxide  —  Bleaching-powder — Cal- 


CONTENTS.  XV 

PAGE 

cium  Carbonate — Applications — Calcium  Sulphate — Calcium 
Phosphates— Calcium  Silicate — Glass— Mortar — Calcium  Sul- 
phide —  Calcium  Nitride  —  Calcium  Carbide  —  Strontium — 
Occurrence  and  Preparation  —  Properties — Compounds  of 
Strontium — Barium — Occurrence  and  Preparation — Properties 
— Barium  Chloride  —  Barium  Hydroxide  —  Barium  Oxide — 
Barium  Peroxide  or  Dioxide  —  Barium  Sulphide  —  Barium 
Nitrate — Barium  Sulphate — Barium  Carbonate — Phosphates  of 
Barium— Reactions  which  are  of  Special  Value  in  Analysis — 
Magnesium  Snb-Group — Gluciuum — Occurrence  and  Prepara- 
tion —  Properties  —  Compounds  of  Glucinum  —  Glucinum 
Chloride — Glucinum  Hydroxide— Glucinum  Sulphate— Glu- 
cinum Carbonate — Weak  Basic  Character  of  Gluciuum — Mag- 
nesium —  Occurrence  —  Preparation — Properties— Applications 
— Compounds  of  Magnesium — Magnesium  Chloride — Mag- 
nesium Oxide — Magnesium  Sulphates — Magnesium  Carbonate 
— Phosphates— Borates— Silicates— Magnesium  Silicide  — Re- 
actions of  Magnesium  Salts  which  are  of  Special  Value  in 
Chemical  Analysis — Erbium — General, 527 

CHAPTER  XXVII. 

ELEMENTS   OF   FAMILY  III,    GROUP   A  : 
ALUMINIUM — SCANDIUM— YTTKIUM — YTTERBIUM— SAMARIUM— HELIUM. 

General  —  Aluminium  —  Occurrence  —  Preparation  —  Properties — 
Applications  —  Aluminium  Chloride  —  Chloroaluminates,  or 
Double  Chlorides  of  Aluminium  and  analogous  Compounds — 
Aluminium  Hydroxide  —  Aluminates — Aluminium  Oxide — 
Aluminium  Sulphate — Basic  Aluminium  Sulphates — Alums — 
Potassium  Alum,  Potassium-Aluminium  Sulphate — Ammo- 
nium Alum,  Ammonium-Aluminium  Sulphate — Sodium  Alum 
— Aluminium  Silicate  —  Kaoline  —  Clay — Ultramarine  —  Por- 
celain— Earthenware — Reactions  of  Aluminium  Salts  which 
are  of  Special  Value  in  Chemical  Analysis — Other  Members  of 
Family  III,  Group  A — Scandium— Yttrium — Ytterbium— Sa- 
marium, Terbium,  and  Gadolinium — Helium — The  Boron- 
Aluminium  Group  in  General 563 

CHAPTER  XXVIII. 

ELEMENTS   OF   FAMILY  I,    GROUP  B  : 
COPPER — SILVER— GOLD. 

General— Copper — General — Forms  in  which  Copper  occurs  in 
Nature — Metallurgy  of  Copper — Properties — Applications — 
Alloys — Cuprous  Hydride — Cupric  Hydride — Cuprous  Chloride 
— Cupric  Chloride— Cuprous  Iodide— Cuprous  Hydroxide— Cu- 
prous Oxide— Cupric  Hydroxide— Cupric  Oxide— Other  Oxides 
of  Copper — -Cupric  Sulphate — Cupric  Nitrate — Cupric  Arseuite 
— Cupric  Carbonates — Cyanides  of  Copper — Cuprous  Sulpho- 


xvi  CONTENTS. 

PAGE 

cyanate — Cupric  Sulphocyanate— Cuprous  Sulphide— Cupric 
Sulphide — Copper-plating— Reactions  which  are  of  Special 
Value  in  Chemical  Analysis— Silver — General — Forms  in  which 
Silver  occurs  in  Nature — Metallurgy  of  Silver— Properties — 
Allotropic  Forms  of  Silver— Alloys  of  Silver — Argentous 
Chloride— Silver  Chloride,  Argentic  Chloride— Silver  Bromide 
and  Iodide — Application  of  the  Chloride,  Bromide,  and  Iodide 
of  Silver  in  the  Art  of  Photography— Silver  Triazoate— Silver 
Oxide— Other  Oxides  of  Silver— Sulphides  of  Silver— Silver 
Nitrate,  Argentic  Nitrate— Silver  Cyanide — Silver  Sulpho- 
cyanate— Berates  of  Silver— Reactions  which  are  of  Special 
Value  in  Chemical  Analysis— Gold— General— Forms  in  which 
Gold  occurs  in  Nature— Metallurgy  of  Gold— Properties- 
Alloys  of  Gold— Chlorides  of  Gold— Chlorauric  Acid— Cyan- 
auric  Acid— Auric  Hydroxide— Gold  Sulphide, 587 

CHAPTER  XXIX. 

ELEMENTS   OF  FAMILY  II,    GROUP  B  : 
ZINC — CADMIUM — MERCURY. 

General — Zinc— General — Forms  in  which  it  occurs  in  Nature- 
Metallurgy — Properties— Applications— Alloys — Zinc  Chloride 
—Zinc  Hydroxide— Zinc  Oxide— Zinc  Sulphide— Zinc  Sulphate 
— Zinc  Carbonate — Reactions  which  are  of  Special  Value  in 
Chemical  Analysis  —  Cadmium  —  General — Preparation  and 
Properties — Cadmium  Sulphide — Cadmium  Cyanide — Analyti- 
cal Reactions — Mercury — General — Forms  in  which  Mercury 
occurs  in  Nature — Metallurgy  of  Mercury — Properties — Appli- 
cations—Amalgams— Mercurous  Chloride — Mercuric  Chloride, 
or  Corrosive  Sublimate — Mercurous  Iodide — Mercuric  Iodide — 
Mercurous  Oxide  —  Mercuric  Oxide  —  Mercurous  Sulphide — 
Mercuric  Sulphide — Mercuric  Cyanide — Mercurous  Nitrate — 
Mercuric  Nitrate — Compounds  formed  by  Salts  of  Mercury 
with  Ammonia — Reactions  which  are  of  Special  Value  in 
Chemical  Analysis, 616 

ELEMENTS   OF  FAMILY  III,  GROUP   B  : 
GALLIUM— INDIUM — THALLIUM. 

General — Gallium — Compounds  of  Gallium — Indium — Compounds 
of  Indium— Thallium— Compounds  of  Thallium, 635 


CHAPTER  XXX. 

ELEMENTS   OF   FAMILY   IV,    GROUP   B: 
GERMANIUM — TIN — LEAD. 

General — Germanium — Tin — General — Occurrence  —  Metallurgy — 
Properties — Applications — Alloys — Stannous  Chloride — Stan- 
nic Chloride  —  Stannous  Hydroxide  —  Stannic  Hydroxide— 
Metastannic  Acid — Stannous  Oxide — Stannic  Oxide — Stannous 


CONTENTS.  xvii 

PAGE 

Sulphide — Stannic  Sulphide — Stannous  and  Stannic  Salts — 
Reactions  which  are  of  Special  Value  in  Chemical  Analysis — 
Lead— General — Forms  in  which  Lead  occurs  in  Nature — Met- 
allurgy —  Properties  - — Applications  —  Lead  Chloride  —  Lead 
Tetrachloride— Lead  Iodide — Lead  Hydroxide — Oxides  of  Lead 
— Lead  Suboxide — Lead  Oxide — Lead  Sesquioxide — Lead  Per- 
oxide— Red  Lead,  Minium — Lead  Sulphide — Lead  Nitrate — 
Lead  Carbonate—Lead  Sulphate — Reactions  which  are  of 
Special  Value  in  Chemical  Analysis— Lanthanum — Cerium — 
Didymium,  Praseodymium  and  Neodymium, 638 

CHAPTER  XXXI. 

ELEMENTS   OF   FAMILY  VI,    GROUP   A. 
CHROMIUM—  MOLYBDENUM — TUNGSTEN — URANIUM/ 

General— Chromium — General— Forms  in  which  Chromium  occurs 
in  Nature — Preparation  —  Properties  —  Chromous  Chloride — 
Chromic  Chloride — Chromous  Hydroxide — Chromic  Hydroxide 
—  Chromic  Oxide  —  Chromic  Sulphate  —  Chrome-Alums — 
Chromic  Acid  and  the  Chromates — Potassium  Chromate — Po- 
tassium Bichromate— Chromium  Trioxide — Relations  between 
the  Chromates  and  Dichromates— Sodium  Chromate  and  So- 
dium Dichromate  —  Barium  Chromate  —  Lead  Chromate  — 
Chromium  Oxychloride,  Chromyl  Chloride— Reactions  which 
are  of  Special  Value  in  Chemical  Analysis — Molybdenum — 
General — Occurrence  and  Preparation — Properties — Chlorides 
— Oxides — Molybdic  Acid  and  the  Molybdates— Lead  Molyb- 
date — Phospho-inolybdic  Acid — Tungsten — General — Occur- 
rence and  Preparation  —  Properties  —  Chlorides  —  Oxides — 
Tungstic  Acid  and  the  Tuugstates— Silico-tungstic  Acids- 
Uranium— General — Occurrence  and  Preparation — Properties 
— Chlorides — Oxides— Uranous  Salts — Urauyl  Salts— Uranates,  657 

CHAPTER  XXXII. 

ELEMENTS  OF  FAMILY  VII,  GROUP  A  : 
MANGANESE. 

General — Forms  in  which  Manganese  occurs  in  Nature — Prepara- 
tion and  Properties— Manganous  Chloride — General  Remarks 
concerning  the  Oxides — Manganous  Oxide — Mnnganous  Hy- 
droxide —  Manganous-manganic  Oxide  —  Manganic  Oxide — 
Mangenese  Dioxide — Manganites — Weldou's  Process  for  the  .  • 
Regeneration  of  Manganese  Dioxide  in  the  Preparation  of 
Chlorine — Sulphides — Mangauous  Cyanide— Mauganous  Car- 
bonate— Mangauous  Sulphate — Manganic  Sulphate — Manganic 
Acid  and  the  Manganates — Permanganic  Acid  and  the  Per- 
manganates—Potassium Permanganate — Reactions  which  are 
of  Special  Value  in  Chemical  Analysis, 678 


xviii  CONTENTS. 

CHAPTER  XXXIII. 

ELEMENTS   OF  FAMILY  VIII,    SUB-GROUP   A  : 
IRON — COBALT — NICKEL. 

PAGK 

Oeneral— Iron— Introductory — Forms  in  which  Iron  occurs  in 
Nature — Metallurgy— Varieties  .of  Iron— Steel — Properties  of 
Iron— Ferrous  Chloride — Ferric  Chloride— Cyanides— Potas- 
sium Ferrocyauide — Ferrohydrocyaiiic  Acid — Ferric  Ferro- 
cyanide,  or  Prussian  Blue — Potassium  Ferricyauide — Ferri- 
hydrocyauic  Acid — Ferrous  Ferricyanide— Nitroprussiates — 
Ferrous  Hydroxide — Ferrous  Oxide — Ferric  Hydroxide— Fer- 
rous-ferric Oxide — Soluble  Ferric  Hydroxide — Ferric  Oxide — 
Ferrous  Sulphide — Ferric  Sulphide  —  Ferrous  Carbonate  - 
Ferrous  Sulphate— Ferric  Sulphate — Ferrous  Phosphate — Fer- 
ric Acid — Iron  Disulphide— Iron  Carbonyls — Reactions  which 
are  of  Special  Value  in  Chemical  Analysis — Ferrous  Compounds 
— Ferric  Compounds — Cobalt— General — Occurrence  and  Prep- 
aration—Properties— Cobaltous  Chloride — Cobaltous  Hydrox- 
ide— Cobaltous  Oxide— Cobaltic  Hydroxide — Cobalt  Sulphide 
— Cyanides — Smalt — Compounds  of  Ammonia  with  Salts  of 
Cobalt— Nickel— General — Occurrence  and  Preparation — Prop- 
erties—  Alloys — Other  Applications  of  Nickel  —  Nickelous 
Chloride — Nickelous  Hydroxide — Nickelic  Hydroxide — Cyan, 
ides — Reactions  of  Cobalt  and  Nickel  which  are  of  Special 
Value  in  Chemical  Analysis, 691 

CHAPTER  XXXIV. 

ELEMENTS   OF  FAMILY   VIII,    SUB-GROUP  B: 

RUTHENIUM— RHODIUM — PALLADIUM. 

ELEMENTS   OF   FAMILY  VIII,    SUB-GROUP   C  : 

OSMIUM— IRIDIUM— PLATINUM. 

General— The  Platinum  Metals— Metallurgy — Ruthenium — Prop- 
erties— Chlorides — Oxides— Osmium — Preparation — Properties 
—  Chlorides  —  Oxides  —  Rhodium  —  Iridium  —  Preparation — 
Properties  —  Chlorides  —  Oxides  —  Palladium  —  Preparation — 
Properties — Palladium — Hydrogen — Chlorides — Oxides — Plati- 
num—Preparation —Properties — Applications  of  Platinum — 
Alleys  of  Platinum — Chlorides — Chlorplatinic  Acid — Cyanides 
— Hydroxides  and  Oxides — Sulphides — Compounds  with  Am- 
monia— The  Platinum  Bases, 719 

APPENDIX  I. 

CONTAINING   SPECIAL    DIRECTIONS   FOR   LABORATORY   WORK. 

Introduction, 733 

EXPERIMENTS   TO   ACCOMPANY   CHAPTER  I. 

Chemical  Change  caused  by  Heat — Chemical  Changes  can  be 
effected  by  an  Electric  Current — Mechanical  Mixtures  and! 
Chemical  Compounds— Other  Examples  of  Chemical  Action,  .  784 


CONTENTS.  xix 

EXPEKTMENT8  TO  ACCOMPANY  CHAPTEB  H. 

PAGE 

Preparation  of  Oxygen— Measurement  of  the  Volume  of  Gases — 
Determination  of  the  Amount  of  Oxygen  liberated  when  a 
known  Weight  of  Potassium  Chlorate  is  decomposed — Physical 
Properties  of  Oxygen — Chemical  Properties  of  Oxygen — Oxy- 
gen 5s  used  up  in  Combustion — The  Products  of  Combustion 
weigh  more  than  the  Body  burned, 740 

EXPERIMENTS  TO  ACCOMPANY  CHAPTEB  III. 

Preparation  of  Hydrogen — Something  besides  Hydrogen  is  formed 
— Determination  of  the  Amount  of  Hydrogen  evolved  when  a 
Known  Weight  of  Zinc  is  dissolved  in  Sulphuric  Acid — Hydro- 
gen is  purified  by  passing  through  a  Solution  of  Potassium 
Permanganate  —  Hydrogen  passes  readily  through  Porous 
Vessels — Diffusion — Chemical  Properties  of  Hydrogen — Prod- 
uct formed  when  Hydrogen  is  Burned — Reduction,  ....  751 

EXPERIMENTS  TO  ACCOMPANY  CHAPTEB  IV. 

Composition  of  Water— Eudiometric  Experiments — Oxyhydrogen 
Blow-pipe,  .../>. 760 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  V. 

Organic  Substances  contain  Water — Water  of  Crystallization- 
Efflorescent  Salts — Deliquescent  Salts — Purification  of  Water 
by  Distillation 763 

EXPERIMENTS  TO  ACCOMPANY  CHAPTEB  VI. 

Method  of  Dumas — Method  of  Victor  Meyer, 765 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  VH. 

Ozone — Hydrogen  Dioxide, 767 

EXPERIMENTS  TO  ACCOMPANY  CHAPTEB  VIII. 

Preparation  of  Chlorine — Chlorine  decomposes  Water  in  the  Sun- 
light— Chlorine  Hydrate — Formation  of  Hydrochloric  Acid— 
Preparation  of  Hydrochloric  Acid 768 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  IX. 

•Chloric  Acid  and  Potassium  Chlorate— Perchloric  Acid,   ....  772 

EXPERIMENTS   TO  ACCOMPANY  CHAPTER   X. 

Neutralization  of  Acids  and  Bases  ;  Formation  of  Salts — Study  of 
the  Products  formed,  .  .  774 


FOR  CHAPTER  XI., 


776 


EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XII. 

Preparation  of  Bromine— Hydrobromic  Acid— Iodine— Iodine  can 
be  detected  by  Means  of  its  Action  upon  Starch -paste — Action 
of  Sulphuric  Acid  upon  Potassium  Iodide  —  lodic  Acid, 
Hydrofluoric  Acid, 776 


XX  CONTENTS. 

EXPERIMENTS~TO  ACCOMPANY  CHAPTER  XIII. 

PAGE 

Properties  of  Sulphur — Hydrogen  Sulphide, 779> 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XIV. 

Manufacture  of  Sulphuric  Acid — Sulphurous  Acid  and  Sulphur 
Dioxide — Sulphurous  Acid  is  a  Reducing  Agent — Sulphur  Tri- 
oxide,  781 

EXPERIMENTS  TO   ACCOMPANY  CHAPTER  XV. 

Preparation  of  Nitrogen— Analysis  of  Air, 784 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XVI. 

Preparation  and  Properties  of  Ammonia — Ammonia  burns  in  Oxy- 
gen— Ammonia  forms  Ammonium  Salts  with  Acids — Compo- 
sition of  Ammonia — Preparation  and  Properties  of  Nitric  Acid 
— Nitric  Acid  gives  up  Oxygen  readily,  and  is  hence  a  good 
Oxidizing  Agent — Metals  dissolve  in  Nitric  Acid,  forming 
Nitrates— Nitrates  are  decomposed  by  Heat — Nitrates  are  sol- 
uble in  Water — Nitric  Acid  is  reduced  to  Ammonia  by  Nascent 
Hydrogen — Nitrous  Acid — Nitrous  Oxide — Nitric  Oxide — 
Nitrogen  Trioxide— Nitrogen  Peroxide, 78$ 

EXPERIMENTS   TO   ACCOMPANY   CHAPTER   XVII. 

Phosphorus — Phosphorus  abstracts  Oxygen  from  other  Substances 
— Phosphine — Arsenic — Arsine — Marsh's  Test  for  Arsenic — 
Antimony  —  Stibine  —  Bismuth  —  Phosphorus  Trichloride  — 
Phosphorus  Pentachloride, .  79ft 

EXPERIMENTS   TO   ACCOMPANY   CHAPTER  XVIII. 

Phosphoric  Acid — Arsenic  Acid — Reduction  of  Arsenic  Trioxide 
—Sulphides  of  Arsenic — Sulphides  of  Antimony — Oxychlorides 
of  Antimony — Basic  Nitrates  of  Bismuth — Boron, 801 

EXPERIMENTS   TO   ACCOMPANY  CHAPTER  XIX. 

Carbon — Bone-black  Filters — Charcoal  absorbs  Gases  —  Carbon 
combines  with  Oxygen  to  form  Carbon  Dioxide — Carbon  re- 
duces some  Oxides  when  heated  with  them — Hydrocarbons,  .  80£ 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XX. 

Carbon  Dioxide  is  formed  when  a  Carbonate  is  treated  with  an 
Acid— Preparation  and  Properties  of  Carbon  Dioxide — Carbon 
Dioxide  is  given  off  from  the  Lungs — Formation  of  Carbon- 
ates— Preparation  and  Properties  of  Carbon  Monoxide — Carbon 
Monoxide  is  a  Good  Reducing  Agent, 80S 

EXPERIMENTS  TO  ACCOMPANY   CHAPTER   XXI. 

Coal  Gas— Oxygen  burns  in  an  Atmosphere  of  a  Combustible  Gas 
—Kindling  Temperature  of  Gases— The  Blow-pipe  and  its 
Uses — Cyanogen 807 


CONTENTS.  xxi 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXH. 

Silicon— Silicon  Tetrafluoride  and  Fluosilicic  Acid— Silicic  Acid,  .  810 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXIV. 

Chlorides,   Bromides,  and   Iodides— Hydroxides— Sulphates— Re- 
duction of  Sulphates  to  Sulphides — Carbonates, 812 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXV. 

Potassium  Salts — Sodium  Salts 816 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXVI. 

Calcium  Salts— Magnesium  and  its  Salts, 817 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXVII. 

Aluminium  Chloride, a 818 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXVIII. 

Copper  and  its  Salts— Silver  and  its  Salts, 818 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXIX. 

Zinc  and  its  Salts— Mercury  and  its  Salts, 819 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXX. 

Tin  and  its  Compounds — Lead  and  its  Compounds, 819 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXXI. 

Chromic  Acid  and  the  Chromates, 820 

EXPERIMENTS  TO   ACCOMPANY  CHAPTER  XXXII. 

Manganese  and  its  Compounds, 821 

EXPERIMENTS   TO   ACCOMPANY  CHAPTER  XXXIII. 

Iron  and  its  Compounds, 821 

EXPERIMENTS   TO   ACCOMPANY   CHAPTER   XXXIV. 

Platinum, 821 

Conclusion, 821 

APPENDIX  II. 

Note — Atomic  Weights — Melting-points  and  Boiling-points  of  the  . 
Elements — Melting-points,  Boiling-points,  and  Solubilities  of 
Inorganic  Substances— Weights  of  Gases  at  0°  and  760  mm. 
Pressure — Approximate  Composition  of  a  number  of  Alloys— 
Freezing-mixtures— Table  of  Weights  and  Measures— Com- 
parison of  the  Twaddell  Scale  with  the  Baume  and  Gay-Lussac 
Scales, 823 


A  TEXT-BOOK  OF 

INORGANIC    CHEMISTRY. 


CHAPTER  I. 

CHEMICAL  AND  PHYSICAL  CHANGE— EARLIEST  CHEMI- 
CAL KNOWLEDGE— LAW  OF  THE  INDESTRUCTIBILITY 
OF  MATTER— LAW  OF  DEFINITE  PROPORTIONS— LAW 
OF  MULTIPLE  PROPORTIONS— THE  ELEMENTS. 

Matter  and  Energy. — The  sensible  universe  is  made  -up 
of  matter  and  energy.  It  is  difficult  to  give  satisfactory 
definitions  of  either  of  these  terms,  but,  in  a  general  way, 
it  may  be  said  that  matter  is  anything  which  occupies 
space,  and  energy  is  that  which  causes  change  in  matter. 
It  requires  but  little  observation  to  show  that  there  are 
many  kinds  of  matter,  and  apparently  many  kinds  of 
energy.  As  examples  of  the  different  kinds  of  matter  we 
have  the  many  varieties  of  rocks  and  earth,  as  granite, 
limestone,  quartz,  clay,  sand,  etc. ;  the  plants  and  their 
fruits ;  the  substances  which  enter  into  the  composition 
of  animals;  and  innumerable  manufactured  products. 
As  examples  of  the  different  forms  of  energy,  we  have 
heat,  light,  motion,  etc.  Under  the  influence  of  the  forms 
of  energy  the  forms  of  matter  are  constantly  undergoing 
change.  Everywhere  these  changes  are  taking  place. 
Changes  in  position  and  in  temperature  appeal  most 
directly  to  our  senses,  and  are  most  easily  studied.  But 
there  are  many  other  kinds  of  change  which  are  of  the 
highest  importance.  Thus  there  are  electrical  changes, 
manifestations  of-  which  we  see  in  thunder-storms ;  there 
are  magnetic  changes  which  may  be  studied  to  some  ex- 
tent by  means  of  the  magnetic  needle ;  and  there  are, 
further,  what  are  called  chemical  changes  which  affect 
the  composition  of  substances. 

(1) 


3  "    ,        INORGANIC  CHEMISTRY. 

Chemical  Change. — For  the  purpose  of  study  it  is  con- 
venient to  distinguish  between  two  classes  of  changes  in 
matter,  the  difference  between  which  can  best  be  made 
clear  by  means  of  examples.  Consider  the  changes  in- 
cluded under  the  head  of  fire.  We  see  substances  de- 
stroyed by  fire,  as  we  say.  They  disappear  as  far  as  we 
can  determine  by  ordinary  observation.  When  iron  is  ex- 
posed to  the  air  a  serious  change  takes  place.  It  becomes 
covered  with  a  reddish-brown  substance  which  we  call 
rust.  If  the  piece  of  iron  is  comparatively  thin,  and  it  be 
allowed  to  lie  in  the  air  long  enough,  it  is  completely 
changed  to  the  reddish-brown  substance,  and  no  iron  as 
such  is  left.  If  the  juices  from  fruits,  as  from  apples,  be 
allowed  to  stand  in  the  air,  they  undergo  change,  becom- 
ing sour,  and  a  somewhat  similar  change  takes  place  in 
milk.  If  a  spark  be  brought  in  contact  with  gunpowder 
there  is  a  flash  and  the  powder  disappears,  a  dense  cloud 
appearing  in  its  place. 

In  the  changes  referred  to  the  substances  changed  dis- 
appear as  such.  After  the  fire,  the  wood  or  the  coal, 
or  whatever  may  be  burned,  is  no  longer  to  be  found. 
The  rusted  iron  is  no  longer  iron.  The  gunpowder  after 
the  flash  is  no  longer  gunpowder.  Changes  of  this  kind 
in  which  the  substances  disappear  and  something  else  is 
formed  in  their  place  are  known  as  cJiemical  changes. 

Physical  Change. — There  are  many  changes  taking 
place  which  do  not  affect  the  composition  of  substances. 
Iron,  for  example,  may  be  changed  in  many  ways  and 
still  remain  iron.  It  may  become  hotter  or  colder.  Its 
position  may  be  changed,  or,  as  we  say,  it  may  be  moved. 
The  iron  may  be  struck  in  such  a  way  as  to  give  forth 
sound.  It  may  be  made  so  hot  that  it  gives  light. 
When,  for  example,  it  becomes  red-hot,  it  can  be  seen  in 
a  dark  room.  A  piece  of  iron  may  be  changed  further 
by  connecting  it  with  what  is  known  as  a  galvanic  bat- 
tery. A  current  of  electricity  then  passes  through  it,  and 
we  can  easily  recognize  the  difference  between  a  piece  of 
iron  through  which  a  current  of  electricit}r  is  passing  and 
one  through  which  no  current  is  passing.  The  former 
when  brought  into  certain  liquids  will  at  once  change 


PHYSICS  AND   CHEMISTRY.  3 

their  composition,  while  the  latter  will  not  cause  such 
change.  Finally,  when  a  piece  of  iron  is  brought  in  con- 
tact with  loadstone,  it  acquires  new  properties.  It  now 
has  the  power  to  attract  and  hold  to  itself  other  pieces 
of  iron.  In  all  these  cases,  the  iron  is  changed,  but  it  re- 
mains iron.  After  the  moving  iron  comes  to  rest,  it  is 
exactly  the  same  thing  that  it  was  before  it  was  moved. 
After  the  iron  which  is  giving  forth  sound  has  ceased  to 
give  forth  sound,  it  returns  to  its  original  condition.  Let 
the  heated  iron  alone,  and  it  cools  down,  ceasing  soon  to 
give  light,  and  presenting  no  evidence  of  being  warm. 
Remove  the  iron  from  contact  with  the  galvanic  battery, 
and  it  loses  those  properties  which  are  due  to  the  current 
of  electricity.  In  time,  the  iron  which  is  magnetized  by 
contact  with  the  loadstone  loses  its  magnetic  properties. 
It  then  no  longer  has  the  power  to  attract  other  pieces 
of  iron  ;  and  does  not  difler  from  ordinary  iron. 

While  iron  has  been  taken  as  an  example,  other  sub- 
stances undergo  similar  changes.  These  changes  which 
do  not  affect  the  composition  of  the  substances  are  called 
physical  changes. 

Physics  and  Chemistry. — According  to  what  has  been 
said,  we  have  two  classes  of  changes  presented  to  us  for 
study : 

(1)  Those  which  do  not  affect  the  composition  of  sub- 
stances, or  physical  changes. 

(2)  Those  which  do  affect   the    composition   of   sub- 
stances, or  chemical  changes. 

That  branch  of  science  which  has  to  deal  with  physical 
changes  is  known  as  PHYSICS.  And  that  which  has  to 
deal  with  chemical  changes  is  known  as  CHEMISTRY. 

Everything  that  has  to  do  with  motion,  heat,  light, 
sound,  electricity,  and  magnetism,  is  studied  under  the 
head  of  Physics.  Everything  that  has  to  do  with  the 
composition  of  substances  is  studied  under  the  head  of 
Chemistry.  It  is,  however,  impossible  to  study  these  two 
subjects  entirely  independently  of  each  other.  When- 
ever a  chemical  change  takes  place,  it  is  accompanied 
by  physical  changes ;  and  in  order  that  the  former  may 
be  clearly  understood,  a  study  of  the  latter  is  necessary. 


INORGANIC  CHEMISTRY. 

Earliest  Chemical  Knowledge. — Those  substances  which 
are  most  abundant  and  most  widely  distributed  in  nature 
were,  of  course,  the  first  known  and  studied ;  and  the 
same  is  true  of  those  chemical  changes  which  occur  most 
commonly  and  produce  the  most  striking  effects.  Simply 
by  observing  those  things  which  surround  us  and  those 
changes  in  composition  which  take  place  naturally,  a 
considerable  amount  of  chemical  knowledge  might  be 
gained,  and  indeed  the  earliest  knowledge  of  chemistry 
was  acquired  in  this  way.  It  was  not,  however,  until 
men  came  to  experiment  upon  the  substances  which  they 
found  in  nature,  that  knowledge  of  chemical  changes 
made  rapid  progress.  Since  then  an  enormous  amount 
of  knowledge  has  been  gained,  and  every  year  the  stock 
is  increased  by  new  discoveries,  until  the  field  appears 
almost  boundless. 

Alchemy. — One  of  the  first  and  one  of  the  most  power- 
ful incentives  to  experiment  upon  chemical  substances 
was  the  desire  to  transform  ordinary  metals  like  lead 
into  gold.  As  will  be  seen  farther  on,  there  was  no  good 
reason  to  suppose  that  transformations  of  this  kind  could 
not  be  effected,  indeed  there  were  good  reasons  for  sup- 
posing them  possible.  For  many  hundred  years  men 
worked  with  this  object  in  view,  and,  though  they  did 
not  succeed  in  accomplishing  that  which  they  undertook, 
they  did  add  greatly  to  our  knowledge  of  the  action  of 
chemical  substances  upon  one  another,  and  they  laid  the 
foundation  of  what  has  since  become  the  great  science  of 
chemistry.  The  work  done  by  the  alchemists  was  neces- 
sary to  prove  that  the  transformations  of  matter  which 
they  tried  to  effect  cannot  be  effected.  The  problem 
which  they  tried  to  solve  was  strictly  a  chemical  prob- 
lem, and  the  work  they  did  was  chemical  work. 

Chemistry  as  a  Science. — While  the  alchemists  accumu- 
lated a  vast  amount  of  knowledge  of  chemical  facts,  they 
did  not,  strictly  speaking,  build  up  the  science  of  chem- 
istry. It  was  only  when  chemists  came  to  study  the  facts 
in  their  relations  to  one  another,  and  when  they  were  en- 
abled to  trace  connections  between  them  ;  when  they  suc- 
ceeded in  discovering  that  certain  general  truths  hold 


LAVOISIER 8   WORK.  5 

good  for  all  cases  of  chemical  action ;  when,  in  short, 
the  fundamental  laws  of  chemical  action  were  discovered 
— it  was  then  that  knowledge  became  science.  It  is 
impossible  to  say  definitely  when  chemistry  became  a 
science.  From  the  unorganized  state  to  the  organized 
there  was  a  gradual  transition.  But  it  is  certain  that  the 
labors  of  the  French  chemist  Lavoisier  contributed 
largely  to  making  chemistry  what  it  is  to-day,  and  it 
is  common  to  refer  to  his  work  as  the  beginning  of  the 
science. 

Lavoisier's  Work. — What  distinguished  Lavoisier's 
work  most  markedly  from  that  of  his  predecessors  was 
the  way  in  which  he  used  the  balance  for  the  purpose  of 
studying  chemical  changes.  The  balance  had  been  used 
to  a  considerable  extent  by  earlier  workers  and  results 
of  value  had  been  reached  by  them,  but  Lavoisier  suc- 
ceeded by  means  of  it  in  explaining  some  important 
Chemical  phenomena  which  had  long  been  the  subject  of 
^tudy.  His  first  investigation,  the  results  of  which  were 
published  in  1770,  was  on  the  transformation  of  water  into 
earth.  It  was  generally  believed  that  water  is  trans- 
formed into  earth  by  boiling,  because  it  was  a  matter  of 
common  observation  that,  whenever  water  is  boiled  for 
a  time  in  a  glass  vessel,  a  deposit  of  earthy  matter  is 
formed.  In  order  to  decide  whether  such  transformation 
takes  place  or  not,  Lavoisier  boiled  some  water  in  a 
closed  vessel  which  he  weighed  before  and  after  the  boil- 
ing ;  and  he  found  that  the  vessel  decreased  in  weight  by 
a  certain  amount.  He  also  determined  the  weight  of  the 
deposit  formed  in  the  vessel  and  found  that  this  was 
exactly  equal  to  the  loss  in  weight  of  the  vessel.  He 
also  showed  that  there  was  just  as  much  water  after  the 
experiment  as  before.  He  therefore  concluded  that  what 
his  predecessors  had  held  to  be  a  transformation  of  water 
into  earth  was  nothing  but  a  partial  disintegration  of  the 
glass  vessel  caused  by  the  action  of  the  boiling  water. 
What  had  appeared  mysterious  became  clear  and  sim- 
ple. Having  succeeded  so  well  in  this  experiment,  La- 
voisier proceeded  to  study  other  chemical  changes  in  the 
same  way,  and  soon  he  was  able  to  give  a  perfectly  satis- 


6  INORGANIC  CHEMISTRY. 

factory  explanation  of  the  process  of  combustion  which 
for  a  long  time  had  been  a  subject  of  study.  The  ex- 
planation and  the  experiments  which  led  to  it  will  be 
taken  up  later.  Suffice  it  to  say  here  that  the  essential 
feature  of  the  work  consisted  in  the  fact  that  the  sub- 
stances which  entered  into  action  and  those  formed 
during  the  action  were  all  carefully  weighed,  and  it  was 
found,  in  every  case,  that  the  weight  of  the  substances 
formed  was  exactly  equal  to  the  weight  of  the  substances 
which  acted  upon  one  another. 

Law  of  the  Indestructibility  of  Matter. — While,  as  has 
been  stated,  chemists  before  Lavoisier  had  used  the 
balance,  they  do  not  appear  to  have  been  very  strongly 
impressed  by  the  importance  of  the  weight  of  sub- 
stances. They  seem  tacitly  to  have  held  that  matter 
can  be  destroyed.  Lavoisier's  work,  however,  showed 
that  whenever  matter  apparently  disappears,  it  continues 
to  exist  in  some  other  form.  If  it  were  possible  to  an- 
nihilate matter  or  to  call  it  into  being,  it  would  be  of 
little  or  no  value  to  weigh  things.  Innumerable  experi- 
ments which  have  been  performed  since  Lavoisier's  time 
have  confirmed  the  view  that  matter  is  indestructible. 
The  first  fundamental  law  bearing  upon  the  changes  in 
composition  which  the  different  forms  of  matter  undergo 
is  the  law  of  the  indestructibility  of  matter.  While,  if  we 
think  of  it,  we  can  scarcely  conceive  that  this  great  law 
should  not  be  true,  we  must  not  forget  that  the  only 
way  in  which  its  truth  could  be  established  was  by  ex- 
periment. The  law  may  be  stated  thus : 

Whenever  a  change  in  the  composition  of  substances  takes 
pt,ace  the  amount  of  matter  after  the  change  is  the  same  as 
before  the  change. 

According  to  this,  and  assuming  that  the  law  has  al- 
ways held  good,  it  follows  that  the  amount  of  matter  in 
the  universe  is  the  same  to-day  as  it  has  been  from  the 
beginning.  Transformations  are  constantly  taking  place, 
but  these  involve  no  increase  nor  decrease  in  the  total 
amount  of  matter. 

Conservation  of  Energy. — Just  as  matter  is  neither  cre- 
ated nor  destroyed,  so  it  has  been  made  probable  that 


COMPOSITION  OF  MATTER.  7 

the  total  amount  of  energy  is  unchangeable.  One  of  the 
greatest  discoveries  in  science  was  the  recognition  of  the 
fact  that  one  form  of  energy  can  be  transformed  into 
others,  and  that  in  these  transformations  nothing  is  lost. 
We  now  know  that  for  a  certain  amount  of  heat  we  can  get 
a  certain  amount  of  motion,  and  that  for  a  certain  amount 
of  motion  we  can  get  a  certain  amount  of  heat.  We  know 
that  a  similar  definite  relation  exists  between  heat  and 
electrical  energy,  and  between  these  and  chemical  energy. 
We  know,  for  example,  that  a  definite  amount  of  heat 
can  be  obtained  by  burning  a  definite  amount  of  a  given 
substance,  and  we  know  also  that  with  a  definite  amount 
of  heat  we  can  produce  a  definite  amount  of  chemical 
change.  Modern  investigation  has  shown  that  all  the 
different  forms  of  energy  are  convertible  one  into  the 
other  without  loss.  This  great  fact  is  generally  spoken 
of  as  the  law  of  the  conservation  of  energy.  Trans- 
formations of  energy  are  taking  place  constantly,  as 
transformations  of  matter  are,  but  the  total  amount  in 
each  case  remains  the  same. 

Early  Views  regarding  the  Composition  of  Matter.— 
The  fact  that  first  impresses  one  in  studying  the  various 
forms  of  matter  found  in  the  earth  is  their  great  variety. 
We  find  an  almost  infinite  number  of  kinds  of  matter, 
and  the  question  at  once  suggests  itself,  of  what  are 
these  things  composed?  This  question  has  long  been 
asked,  and  it  will  be  long  before  an  entirely  satisfactory 
answer  is  reached.  Still,  much  more  is  now  known  in 
regard  to  the  subject  than  was  known  in  past  ages,  and 
some  progress  is  constantly  being  made  towards  a  solu- 
tion of  the  problem.  At  first,  men  attempted  to  answer 
the  question,  as  they  attempted  to  answer  all  important 
questions,  by  what  is  known  as  the  speculative  method ; 
that  is  to  say,  they  took  the  facts,  as  far  as  they  knew 
them,  into  consideration,  and  they  endeavored  by  purely 
mental  processes  to  furnish  an  explanation.  They  were 
on  the  whole  much  bolder  in  the  use  of  the  imagination 
than  the  scientific  men  of  the  present  are,  or  rather  they 
do  not  appear  to  have  had  as  great  respect  for  facts  as 
men  now  have,  and,  as  a  consequence,  we  find  that  some 


8  INORGANIC  CHEMISTRY. 

extremely  curious  speculations  were  indulged  in.  One 
of  the  most  prominent  views  in  regard  to  the  composi- 
tion of  matter  was  that  put  forward  by  Aristotle.  Ac- 
cording to  this  view,  all  forms  of  matter  are  made  up  of 
four  elements,  earth,  air,  fire,  and  water ;  and  the  vari- 
ous forms  differ  from  one  another  in  the  proportions  of 
these  elements  contained  in  them.  Aristotle  evidently 
had  in  mind  the  fundamental  properties  of  the  four  ele- 
ments, rather  than  the  elements  themselves,  and  his  idea 
was  that  these  fundamental  properties  are  found  in  dif- 
ferent proportions.  Instead  of  meaning  that  water  as 
such  was  contained  in  substances,  he  meant  that  the 
properties  of  cold  and  moisture  were  met  with  in  sub- 
stances, and  so  on  for  the  other  elements ;  fire  repre- 
senting heat  and  dryness ;  earth,  cold  and  dryness ;  and 
air,  heat  and  moisture.  Afterwards  it  was  pointed  out 
that  besides  the  four  fundamental  properties  represented 
by  the  four  elements,  each  substance  has  a  special  prop- 
erty of  its  own  which  distinguishes  it  from  all  others. 
This  was  called  the  quinta  essentia,  or  fifth  essence,  from 
which  our  modern  word  quintessence  is  derived.  The 
four  or  five  elements  of  the  older  philosophers  were,  as 
will  be  seen,  imaginary  things.  They  represented  ideas 
rather  than  tangible  substances. 

Elements. — As  experimenting  upon  chemical  substances 
advanced  further  and  further,  the  fact  impressed  itself 
more  and  more  strongly  upon  investigators,  that,  of  the 
large  number  of  substances  known,  some  can  be  con- 
verted into  simpler  ones  by  chemical  action  and  some 
cannot.  In  other  words,  some  substances  like  water 
can  be  broken  down  by  various  methods  into  two  or 
more  others  of  different  properties,  and  these  when 
brought  together  again  under  proper  conditions  form 
the  original  substance.  In  the  case  of  water,  the  action 
of  an  electric  current  breaks  it  down  or  decomposes  it, 
and  two  gases,  hydrogen  and  oxygen,  are  formed  from  it. 
Elaborate  experiments  have  shown  that  the  weight  of 
water  decomposed  is  exactly  equal  to  the  weight  of  the 
hydrogen  plus  that  of  the  oxygen  obtained,  and  that 
when  the  hydrogen  and  oxygen  are  brought  together 


ELEMENTS.  9 

again  under  proper  conditions  exactly  as  much  water  is 
formed  as  was  originally  decomposed.  It  appears,  there- 
fore, that  water  consists  of  at  least  two  simpler  sub- 
stances. A  similar  conclusion  is  reached  by  a  study  of 
by  far  the  largest  number  of  the  substances  with  which 
we  have  to  deal.  On  the  other  hand,  no  treatment  to 
which  hydrogen  and  oxygen  have  been  subjected  has,  as 
yet,  effected  their  decomposition.  They  can  be  made  to 
combine  with  other  substances,  as,  for  example,  with  ^ach 
other,  and  thus  form  more  complex  substances,  but  noth- 
ing simpler  than  hydrogen  has  ever  been  obtained  from 
hydrogen,  and  nothing  simpler  than  oxygen  has  ever  been 
obtained  from  oxygen.  Whether  the  decomposition  of 
these  substances  will  ever  be  effected  is  a  question  which 
cannot  be  answered.  All  that  we  know  is  that  at  present 
they  cannot  be  decomposed.  We  therefore  speak  of 
them  as  elements,  meaning  by  the  term,  that,  with  the 
means  now  at  the  disposal  of  chemists,  it  is  impossible  to 
get  simpler  substances  from  them.  There  are  at  present 
about  seventy  substances  known  which  are  called  ele- 
ments for  the  same  reasons  that  hydrogen  and  oxygen 
are  called  elements.  It  is  quite  possible  that  the  num- 
ber may  be  increased  in  the  future,  and  it  is  also  quite 
possible  that  the  number  may  be  decreased.  New  ele- 
ments will  in  all  probability  be  discovered,  and  prob- 
ably some  of  the  substances  now  included  in  the  list  of 
elements  may  eventually  be  shown  to  be  capable  of 
decomposition. 

The  view  at  present  held  in  regard  to  the  forms  of 
matter  which  go  to  make  up  that  part  of  the  universe 
which  comes  under  our  observation  is  that  they  are  all 
composed  of  the  seventy  elementary  substances.  Many 
of  them,  like  water,  are  composed  of  only  two  elements ; 
others  of  three ;  and  still  others  of  four,  five,  six,  and 
more  ;  but  most  of  them  are  comparatively  simple,  and 
rarely  does  any  one  contain  more  than  four  or  five  ele- 
ments. Of  the  seventy  elements  known,  only  about 
twelve  enter  into  the  composition  of  most  things  with 
which  we  commonly  have  to  deal.  The  others  occur  in 
.relatively  small  quantity. 


10  INORGANIC  CHEMISTRY. 

Chemical  Action. — In  the  last  paragraph  it  was  stated 
that  most  substances  can  be  decomposed,  and  that  under 
proper  conditions  the  elements  combine.  We  must  now 
inquire  more  carefully  into  the  meaning  of  these  expres- 
sions. Among  the  elements  are  the  well-known  sub- 
stances lead,  iron,  and  sulphur.  If  some  finely  divided 
iron  is  brought  in  contact  with  sulphur,  apparently  no 
action  takes  place.  If  the  two  are  put  in  a  mortar  and 
mixed  no  matter  how  thoroughly,  there  is  no  evidence 
of  action.  The  mixture  has,  to  be  sure,  a  different  ap- 
pearance from  that  of  either  constituent,  but  still  both 
substances  are  present,  and  can  be  recognized  by  various 
methods.  If,  for  example,  a  little  of  the  mixture  is  ex- 
amined with  the  aid  of  the  microscope,  particles  of  iron 
and  of  sulphur  will  be  recognized  lying  side  by  side.  If, 
further,  the  mixture  is  treated  with  the  liquid,  carbon 
disulphide,  which  has  the  power  to  dissolve  the  sulphur 
but  not  the  iron,  the  sulphur  will  be  dissolved  while  the 
iron  will  be  left  unchanged.  Finally,  if  a  dry  magnet  is 
introduced  into  the  mixture,  the  iron  will  adhere  to  it, 
and  by  careful  manipulation  the  two  constituents  can  be 
separated.  These  facts  furnish  evidence  that  both  iron 
and  sulphur  are  present  in  the  mixture  in  unchanged 
condition,  just  as  sugar  and  sand  are  present  in  a  mixture 
of  these  two  substances.  If  now  the  mixture  of  sulphur 
and  iron  is  heated  in  a  dry  test-tube,  marked  changes 
will  take  place,  and  there  will  be  formed  a  black  sub- 
stance entirely  different  from  either  of  the  elements  em- 
ployed in  the  experiment.  Carbon  disulphide  can  no 
longer  extract  sulphur  from  it.  The  magnet  can  no 
longer  pick  out  the  iron,  and  under  the  microscope  one 
homogeneous  substance  is  seen  instead  of  the  two  ele- 
mentary substances.  If  the  experiment  is  performed 
with  proper  precautions,  the  amount  of  matter  after  the 
action  will  be  found  to  be  exactly  the  same  as  before  the 
action.  A  serious  change  has  taken  place,  but  no  change 
in  the  amount  of  matter.  The  act  is  one  of  chemical  com- 
bination, and  the  substance  formed  is  called  a  chemical 
compound.  A  few  other  examples  will  aid  in  making  the 
conception  of  chemical  combination  clear.  When  a  bit 


CHEMICAL  ACTION  AND  AFFINITY.  11 

of  phosphorus  is  brought  in  contact  with  a  little  iodine 
action  takes  place  at  once ;  the  two  elements  combine, 
losing  their  own  characteristic  properties  and  forming  a 
compound  with  properties  quite  different  from  those  of 
the  constituents.  When  the  gases  hydrogen  and  oxygen 
are  brought  together  and  a  spark  is  passed  through  the 
mixture  an  explosion  occurs,  and,  in  place  of  the  gases,  the 
liquid,  water,  is  formed.  When  sulphur  burns  in  the  air 
the  product  formed  is  a  pungent  gas.  It  has  been  shown 
that  the  act  consists  in  the  combination  of  the  sulphur 
with  the  gas,  oxygen,  which  is  contained  in  the  air.  All 
these  cases  are  examples  of  chemical  combination.  But 
chemical  action  may  be  of  the  opposite  kind,  that  is  to 
say,  instead  of  being  combination,  it  may  be  decomposition. 
Thus,  water  which  is  formed  by  the  chemical  combina- 
tion of  hydrogen  and  oxygen  may,  by  proper  methods, 
be  decomposed  into  the  same  elements.  We  may  con- 
veniently think  of  that  which  causes  elements  to  combine 
as  an  attractive  force  exerted  between  the  elements. 
Now,  when  some  power  which  can  overcome  this  attrac- 
tion is  brought  to  bear  upon  a  compound,  decomposition 
takes  place,  and  the  elements  are,  as  we  say,  set  free. 
When,  for  example,  an  electric  current  is  passed  through 
water,  the  power  which  holds  together  the  hydrogen  and 
oxygen  is  overcome  and  bubbles  of  the  one  gas  rise  from 
one  pole  of  the  battery  and  bubbles  of  the  other  gas  rise 
from  the  other  pole.  This  is  a  simple  example  of  chemi- 
cal decomposition.  Again,  when  the  substance  known  as 
red  oxide  of  mercury  or  mercuric  oxide  is  heated  to  a 
sufficiently  high  temperature  a  colorless  gas  is  given  off 
from  it,  and  globules. of  mercury  are  formed  at  the  same 
time.  The  gas,  as  will  be  shown  later,  is  oxygen,  so  that 
from  the  red  oxide  of  mercury,  which  is  a  chemical  com- 
pound of  mercury  and  oxygen,  we  get,  by  heating,  the 
two  elements  in  the  free  state.  In  this  case,  heat  over- 
comes the  chemical  attraction  which,  in  the  compound, 
holds  the  elements  together. 

Chemical  Affinity. — It  is  evident  from  what  has  already 
been  said  that  there  is  some  power  which  can  hold  sub- 
stances together  in  a  very  intimate  way,  so  intimate  that 


12  INORGANIC  CHEMISTRY. 

we  cannot  recognize  them  by  ordinary  means.  We  do 
not  know  what  causes  the  sulphur  and  iron  to  combine, 
but  we  do  know  that  they  combine.  Similarly,  we  do  not 
know  what  causes  a  stone  thrown  in  the  air  to  fall  back 
again,  but  we  know  that  it  falls  back.  It  is  true  we  may 
say  that  the  cause  of  the  falling  of  the  stone  is  the  attrac- 
tion of  gravitation,  but  this  does  not  give  us  any  real  in- 
formation, for,  if  we  ask  what  the  attraction  of  gravitation 
is,  we  can  only  answer  that  it  is  that  which  causes  all 
bodies  to  attract  one  another.  We  may  also  say,  and  do 
say,  that  the  cause  of  chemical  combination  is  chemical 
affinity.  But  in  so  doing  we  only  give  a  name  to  something 
about  which  we  know  nothing  except  the  effects  it  pro- 
duces. All  the  chemical  changes  which  are  taking  place 
around  us  may,  then,  be  referred  to  the  operation  of  chemi- 
cal affinity.  If  this  power  should  cease  to  operate,  what 
would  be  the  result?  Nature  would  be  infinitely  less 
complex  than  it  now  is.  All  complex  substances  wrould 
be  resolved  into  the  elements  of  which  they  are  com- 
posed, and,  as  far  as  we  know,  there  would  be  only  about 
seventy  different  kinds  of  substances.  All  living  things 
would  cease  to  exist,  and  in  their  place  there  would  be 
three  invisible  gases,  and  something  very  much  like  char- 
coal. Mountains  would  crumble  to  pieces,  and  all  water 
would  disappear  giving  two  invisible  gases.  The  pro- 
cesses of  life  in  its  many  forms  would  be  impossible,  as, 
however  subtle  that  which  we  cpll  life  may  be,  we  cannot 
imagine  it  to  exist  without  the  existence  of  certain  com- 
plex forms  of  matter ;  and,  as  regards  the  life  process 
of  animals  and  plants,  most  complex  chemical  changes 
are  constantly  taking  place  within  them,  and  these 
changes  are  essential  to  the  continuance  of  life. 

Chemical,  Compounds  and  Mechanical  Mixtures. — The 
substances  formed  by  chemical  combination  of  the  ele- 
ments are  called  chemical  compounds.  Most  substances 
found  in  nature  are  made  up  of  several  others.  Wood, 
for  example,  is  very  complex,  containing  a  large  number 
of  distinct  chemical  compounds  intimately  mixed  together. 
Some  of  these  can  be  isolated,  but  it  is  impossible  to 
isolate  them  all  wHh  the*  means  at  present  at  our  com- 


CHEMICAL  COMPOUNDS  AND  MECHANICAL  MIXTURES.  13 

mand.  Most  of  the  rocks  met  with  are  also  quite  com- 
plex, and  it  is  difficult  to  isolate  the  constituents.  If  we 
look  at  a  piece  of  coarse-grained  granite,  we  see  plainly 
enough  that  it  contains  different  things  mixed  together, 
and  if  it  be  broken  up  we  can  pick  out  pieces  of  different 
substances  from  the  mass.  If  we  now  examine  a  piece 
of  each  of  the  different  substances  thus  picked  out  of  the 
granite,  it  appears  to  be  homogeneous,  i.e.  we  cannot 
recognize  the  presence  of  more  than  one  kind  of  thing  in 
-any  one  piece.  If  the  piece  is  carefully  selected  it  may 
be  powdered  finely  in  an  agate  mortar,  and  some  of  the 
powder  examined  with  a  microscope  without  the  presence 
of  more  than  one  substance  being  recognized.  We  are 
&ble  to  isolate  three  substances  from  granite  by  simply 
breaking  it  up  and  picking  out  the  pieces  of  different 
kinds.  We  might  therefore  conclude  that  granite  con- 
sists of  three  substances.  This  is  true,  but  it  is  not  the 
whole  truth.  For  it  is  possible  by  proper  means  to  get 
simpler  substances  from  each  of  the  three  already  sep- 
arated. This  is,  however,  a  much  more  difficult  process 
than  the  separation  first  accomplished.  To  effect  the  sep- 
aration of  each  of  the  three  constituents  of  granite  into 
its  elements  requires  more  powerful  means.  Substances 
must  be  brought  in  contact  with  them  which  act  upon 
them,  changing  their  composition,  i.e.  act  chemically 
upon  them,  and  high  heat  must  be  used  to  aid  the  action. 
By  skilful  work  it  is,  however,  possible  to  separate  the 
three  components  of  granite  into  their  elements. 

From  the  above  it  is  evident  that  substances  may  be 
united  in  different  ways.  They  may  be  so  united  that  it  is 
a  simple  thing  to  separate  them  by  mechanical  processes. 
Or  they  may  be  so  united  that  it  is  impossible  to  separate 
them  by  mechanical  processes.  By  a  mechanical  process 
is  meant  any  process  which  does  not  involve  the  use  of 
heat,  electricity,  or  chemical  change.  Thus,  the  mechan- 
ical process  made  use  of  in  the  case  of  granite  consisted 
in  picking  out  the  pieces.  The  separation  of  the  parti- 
cles of  different  sizes  by  means  of  a  sieve  is  a  mechanical 
process.  The  separation  of  two  liquids  which  do  not  mix 
with  each  other  is  a  mechanical  process.  Complex  sub- 


14  INORGANIC  CHEMISTRY. 

stances  which  may  be  separated  into  their  components  by 
purely  mechanical  processes  are  called  mechanical  mix- 
tures. Thus  granite  is  a  mechanical  mixture  of  three 
chemical  compounds.  Similarly,  most  natural  substances 
are  more  or  less  complex  mixtures  of  chemical  com- 
pounds, or,  much  more  rarely,  of  elements.  Air,  for  ex- 
ample, is  a  mechanical  mixture  consisting  mainly  of  the 
two  elements  nitrogen  and  oxygen.  It  is  not  always  an 
easy  matter  to  distinguish  between  mechanical  mixtures 
and  chemical  compounds,  as  there  are  mixtures  which  it  is 
extremely  difficult  to  subdivide  into  their  components,  and 
there  are,  on  the  other  hand,  chemical  compounds  which 
are  extremely  unstable.  Generally,  however,  the  differ- 
ence is  recognized  without  serious  difficulty. 

Qualitative  and  Quantitative  Study  of  Chemical  Changes. 
— In  general  there  are  two  ways  in  which  chemical 
changes  may  be  studied.  Substances  may  be  brought 
together  under  a  variety  of  conditions  and,  if  action  takes 
place,  the  properties  of  the  product  or  products  may  then 
be  studied  and  compared  with  those  of  the  substances 
brought  together.  In  the  early  periods  of  the  history  of 
chemistry  the  study  was  almost  wholly  of  this  kind.  This 
is  called  qualitative  study.  But  we  may  go  farther  than 
this,  and  take  into  consideration  the  weights  or  masses 
of  the  substances  we  are  dealing  with.  We  should  then 
be  studying  the  changes  quantitatively.  We  have  already 
seen  that  by  means  of  the  quantitative  method  Lavoisier 
placed  the  law  of  the  %  indestructibility  of  matter  upon  a 
firm  basis,  and  that  he  also  succeeded  by  the  use  of  this 
method  in  explaining  a  number  of  important  chemical 
changes,  particularly  combustion.  By  further  use  of  this 
method  other  laws  of  the  highest  importance  to  the  science 
of  chemistry  were  soon  brought  to  light. 

Law  of  Definite  Proportions. — The  fact  that  sulphur  and 
iron  combine  chemically  when  a  mixture  of  the  two  is 
heated  has  been  referred  to.  The  question  whether  they 
combine  in  all  proportions  is  one  which  can  be  answered 
only  by  a  quantitative  study  of  the  process.  If  the  pro- 
cess were  to  be  studied  for  the  first  time  the  method  of 
procedure  would  be  this  :  We  should  mix  the  elements 


LAW  OF  DEFINITE  AND  OF  MULTIPLE  PROPORTIONS.    15 

in  different  proportions  and,  after  the  action,  we  should 
determine  whether  any  of  either  of  the  elements  is  left  in 
the  uncombined  state  ;  and,  further,  by  decomposing  the 
product,  we  should  determine  whether  it  always  contains 
the  elements  in  the  same  proportions.  The  problem,  in  this 
case,  is  by  no  means  a  simple  one,  but  it  has  been  repeat- 
edly worked  over  with  the  greatest  possible  care,  and,  as 
the  result  of  the  work,  the  conclusion  is  justified  that  the 
product  always  contains  the  elements  in  exactly  the  same 
proportions.  Similar  work  has  been  done  for  most  other 
chemical  compounds  known,  and  the  general  conclusion 
known  as  the  law  of  definite  proportions  has  been  drawn. 
This  law  may  be  stated  thus : 

A  chemical  compound  always  contains  the  same  constitu- 
ents in  the  same  proportion  by  weight. 

The  truth  of  this  general  statement  or  law  has  not  al- 
ways been  acknowledged  by  chemists.  At  the  beginning 
of  this  century  a  celebrated  discussion  on  the  subject  took 
place  between  Proust  and  Berthollet.  The  discussion 
led  to  a  great  deal  of  careful  work  which  tended  to  con- 
firm the  law,  and  since  that  time  it  has  not  been  seriously 
doubted.  About  twenty  years  ago  a  Belgian  chemist, 
Stas,  by  a  long  series  of  probably  the  most  painstaking 
and  accurate  chemical  experiments  ever  performed, 
showed  that  in  the  compounds  which  he  worked  with  there 
was  no  variation  in  composition  that  could  be  detected 
by  the  most  refined  methods  of  chemistry.  In  the  pres- 
ent state  of  our  knowledge  it  appears  that  the  law  of 
definite  proportions  is  a  law  in  the  strictest  sense. 

Law  of  Multiple  Proportions. — It  does  not  require  a  very 
extended  study  of  chemical  phenomena  to  show  that  from 
the  same  elements  it  is  possible  in  many  cases  to  get 
more  than  one  product.  Thus  iron  and  sulphur  form 
three  distinct  compounds  with  each  other.  Tin  combines 
with  oxygen  in  two  proportions.  The  elements  potassium, 
chlorine,  and  oxygen  combine  in  four  different  ways,  form- 
ing four  distinct  products.  Nitrogen  and  oxygen  form 
five  products.  In  the  early  part  of  this  century  the  Eng- 
lish chemist  Dalton  by  a  study  of  cases  like  those  men- 
tioned was  led  to  the  discovery  of  another  great  law  of 


16  INORGANIC  CHEMISTRY. 

chemistry  known  as  the  law  of  multiple  proportions* 
Many  substances  had  been  analyzed  before  his  time,  and 
the  percentages  of  the  constituents  determined  with  a  fair 
degree  of  accuracy.  He  examined,  first,  two  gases,  both 
of  which  consist  of  carbon  and  hydrogen.  He  determined 
the  percentages  of  the  constituents,  and  found  them  to- 
be  as  follows : 

Olefiant  gas,  85.7  per  cent,  carbon  and  14.3  per  cent, 
hydrogen. 

Marsh  gas,  75.0  per  cent,  carbon  and  25.0  per  cent,  hy- 
drogen. 

On  comparing  these  numbers,  he  found  that  the  ratio  of 
carbon  to  hydrogen  in  olefiant  gas  is  as  6  to  1 ;  whereas 
in  marsh  gas  it  is  as  3  to  1  or  6  to  2.  The  mass  of  hy- 
drogen, combined  with  a  given  mass  of  carbon,  is  exactly 
twice  as  great  in  the  one  case  as  in  the  other. 

There  are,  further,  two  compounds  of  carbon  and  oxy- 
gen, and  in  analyzing  these  the  following  figures  were 
obtained : 

Carbon  monoxide,  42.86  per  cent,  carbon  and  57.14  per 
cent,  oxygen. 

Carbon  dioxide,  27.27  per  cent,  carbon  and  72.73  per 
cent,  oxygen. 

But   42.86  :  57.14  ::  6  :  8   and    27.27  :  72.73  ::  6  :  16. 

The  mass  of  oxygen  combined  with  a  given  mass  of 
carbon  in  carbon  dioxide  is  exactly  twice  as  great  as  the 
mass  of  oxygen  combined  with  the  same  mass  of  carbon  in 
carbon  monoxide.  These  facts  and  other  similar  ones  led 
to  the  discovery  of  the  law  of  multiple  proportions,  which 
may  be  stated  thus: 

If  two  dements  A  and  B  form  several  compounds  with 
each  other,  and  we  consider  any  fixed  mass  of  A,  then  the 
different  masses  of  B  which  combine  with  the  fixed  mass  of 
A  bear  a  simple  ratio  to  one  another. 

By  way  of  further  illustration  we  may  take  the  three 
compounds  which  iron  forms  with  sulphur.  In  one  of 
these,  approximately  7  parts  of  iron  are  in  combination 
with  4  parts  of  sulphur  ;  in  a  second,  7  parts  of  iron  are  in 
combination  with  6  parts  of  sulphur ;  and  in  the  third,  7 
of  iron  are  in  combination  with  8  of  sulphur.  The  figures 


COMBINING    WEIGHTS  OF  THE  ELEMENTS.          17 

'4,  6,  and  8  bear  a  simple  ratio  to  one  another  which  is 
2:3:4.  The  five  compounds  of  nitrogen  and  oxygen 
contain  7  parts  of  nitrogen  combined  with  8,  16,  24,  32, 
and  40  parts  of  oxygen  respectively.  The  figures  repre- 
senting the  parts  by  weight  of  oxygen  combined  with  7 
parts  by  weight  of  nitrogen  are  in  the  ratio  1:2:3:4:5. 
In  the  compounds  formed  by  the  elements  chlorine, 
potassium,  and  oxygen  the  proportions  by  weight  are 
represented  in  the  following  table  : 

Chlorine.  Potassium.  Oxygen. 
35.18                        38.82  15.88 

35.18  38.82  31.76 

35.18  38.82  47.64 

35.18  38.82  63.52 

It  will  be  observed  that  the  ratio  between  the  chlorine 
and  potassium  remains  constant,  but  that  the  mass  of 
oxygen  varies  regularly  from  15.88  to  63.52 ;  the  masses 
bearing  to  one  another  the  simple  ratio  1:2:3:4. 

The  law  of  multiple  proportions  like  the  law  of  defi- 
nite proportions  is  simply  a  statement  in  accordance 
with  what  has  been  found  true  by  experiment.  Although 
discovered  by  Dalton  at  the  beginning  of  this  century 
and  put  forward  upon  what  appears  now  to  be  only  a 
slight  basis  of  facts,  all  work  since  that  time  has  con- 
firmed it,  and  it  forms  to-day  one  of  the  corner-stones  of 
the  science  of  chemistry. 

Combining  Weights  of  the  Elements. — A  careful  study 
of  the  figures  representing  the  composition  of  chemical 
compounds  reveals  a  remarkable  fact  regarding  the  rela- 
tive quantities  of  one  and  the  same  element  which  enter 
into  combination  with  different  elements.  The  propor- 
tions by  weight  in  which  some  of  the  elements  combine 
chemically  with  one  another  are  stated  in  the  following 
table  : 

1  part  Hydrogen  combines  with  35.18  parts  Chlorine. 


1     "  "  "  "     79.34 

1     "  "  "  "  125.89 

35.18  parts  Chlorine  combine  "     38.82 

79.34     "     Bromine        "         "     38.82 

125.89     "     Iodine  "         "     38.82 


Bromine. 

Iodine. 

Potassium. 


18 


INORGANIC  CHEMISTRY. 


15.88  parts  Oxygen  combine  with  64.91  parts  Zinc. 


15.88 

" 

15.88 

" 

15.88 

" 

64.91 

Zinc 

34.10 
39.76 

Magnesium 
Calcium 

136.39 

Barium 

24.10 
b9.76 

Magnesium. 
Calcium. 

136.39 

Barium. 

31.83 
31.83 

Sulphur. 

31.83 

" 

31.83 

« 

It  will  be  seen  that  the  figures  which  express  the  rela- 
tive quantities  of  chlorine,  bromine,  and  iodine  that 
combine  with  1  part  of  hydrogen  also  express  the  rela- 
tive quantities  of  these  elements  that  combine  with 
38.82  parts  of  potassium.  So  also  the  figures  which  ex- 
press the  relative  quantities  of  zinc,  magnesium,  calcium, 
and  barium  that  combine  with  15.88  parts  of  oxygen 
also  express  the  relative  quantities  of  these  elements 
that  combine  with  31.83  parts  of  sulphur.  Now,  an 
examination  of  all  compounds  known  has  shown  that 
hydrogen  enters  into  combination  with  the  other  elements 
in  the  smallest  proportions  ;  it  is  therefore  taken  as  unity 
in  stating  the  relative  quantities  of  the  other  elements 
which  enter  into  combination.  The  weight  of  another 
element  that  combines  with  1  part  by  weight  of  hydro- 
gen may  be  called  its  combining  iveight.  Thus,  according 
to  the  above,  the  combining  weights  of  chlorine,  bromine, 
and  iodine  are  respectively  35.18,  79.34,  and  125.89. 
Similarly  38.82  is  the  combining  weight  of  potassium, 
as  it  expresses  the  weight  of  potassium  that  combines 
with  the  above  quantities  of  chlorine,  bromine,  and 
iodine.  Thus  for  every  element  a  number  can  be  se- 
lected, such  that  the  proportions  by  weight  in  which  the 
element  enters  into  combination  with  others  can  be  con- 
veniently expressed  by  the  number  or  by  a  simple  multiple 
of  it.  These  numbers  are  the  combining  weights. 

It  is  not  by  any  means  an  easy  matter  to  determine 
which  numbers  are  most  convenient  for  all  circumstances  ; 
and  if  the  selection  is  to  be  determined  solely  by  con- 
venience, there  may  be  differences  of  opinion  as  to  what 
is  most  convenient.  We  shall  see  a  little  later  that, 
while  the  numbers  primarily  express  the  combining 


SYMBOLS  AND  ATOMIC  WEIGHTS  OF  THE  ELEMENTS.  19 


weights  and  nothing  else,  and  are  based  solely  upon  a 
study  of  the  composition  of  chemical  compounds,  they 
have  come  to  have  a  deeper  significance  which  will  be 
explained  in  due  time.  They  are  now  called  atomic 
iveights  because  for  strong  reasons  they  are  believed  to 
express  the  relative  weights  or  masses  of  the  minute  in- 
divisible particles  or  atoms  of  which  the  various  kinds  of 
matter  are  assumed  to  be  made  up.  The  atomic  theory, 
as  it  is  called,  will  be  treated  of  farther  on,  and  the  re- 
lation which  exists  between  the  theory  and  the  figures 
called  the  atomic  weights  will  be  discussed  at  some 
length.  For  the  present  it  will  be  best  to  use  the  figures 
as  expressing  the  combining  weights,  and  as  being  en- 
tirely independent  of  any  speculations  regarding  the  con- 
stitution of  matter  and  the  existence  of  atoms. 

The  Elements,  their  Symbols  and  Atomic  Weights. — It 
has  already  been  stated  that  there  are  about  seventy 
elementary  substances  known,  but  that  of  these  only  a 
small  number  enter  into  the  composition  of  common 
things  to  any  great  extent.  It  has  been  calculated  that 
the  solid  crust  of  the  earth  is  made  up  approximately  as 
represented  in  the  subjoined  table : 


Oxygen 47.13$ 

Silicon 27.89$ 

Aluminium 8.13$ 

Iron...  4.71$ 


Calcium 3.53$ 

Magnesium 2. 64$ 

Sodium 2.68$ 

Potassium .  2.35$ 


While  oxygen  forms  a  large  proportion  of  the  solid 
crust  of  the  earth,  it  forms  a  still  larger  proportion  (eight- 
ninths)  of  water,  and  about  one-fifth  of  the  air.  Carbon 
is  the  principal  element  which  enters  into  the  structure 
of  living  things,  while  hydrogen,  oxygen,  and  nitrogen 
also  are  essential  constituents  of  animals  and  plants. 
Nitrogen  forms  about  four-fifths  of  the  air. 

In  representing  the  results  of  chemical  action,  it  is  con- 
venient to  use  abbreviations  for  the  names  of  the  elements 
and  compounds.  Thus,  instead  of  oxygen  we  may  write 
simply  O  ;  for  hydrogen,  H  ;  for  nitrogen,  N ;  etc.  These 
symbols  are  used  in  expressing  what  takes  place  when  sub- 


20  INORGANIC  CHEMISTRY. 

stances  act  upon  one  another.  Very  frequently  the  first 
letter  of  the  name  is  used  as  the  symbol.  If  the  names- 
of  two  or  more  elements  begin  with  the  same  letter, 
this  letter  is  used,  and  some  other  letter  of  the  name  is 
added.  Thus,  B  is  the  symbol  of  boron,  Ba  of  barium, 
Bi  of  bismuth  ;  C  is  the  symbol  of  carbon,  Ca  of  calcium, 
Cd  of  cadmium,  Ce  of  cerium,  Cl  of  chlorine,  Cr  of 
chromium,  Cs  of  caesium,  Cu  of  copper.  In  some  cases 
the  symbol  is  derived  from  the  Latin  name  of  the  ele- 
ment. Thus,  the  symbol  of  iron  is  Fe,  from  the  Latin 
ferrum;  for  copper  Cu,  from  cuprum;  for  mercury  Hg> 
from  hydrargyrum;  etc. 

The  names  themselves  are  formed  in  a  variety  of  ways. 
Chlorine  is  derived  from  the  Greek  ^/Vc^po?,  which  means- 
yellowish-green,  as  this  is  the  color  of  chlorine.  Bro- 
mine comes  from  /?pc3/*o£,  a  stench,  a  prominent  charac- 
teristic of  bromine  being  its  bad  odor.  Hydrogen  comes, 
from  vdoop,  water,  and  ysreiv,  to  produce,  signifying  that 
it  is  a  constituent  of  water.  Similarly  nitrogen  comes- 
from  virpov,  niter,  and  yeveiv,  to  produce,  nitrogen  be- 
ing one  of  the  constituents  of  niter.  Potassium  is  an 
element  found  in  potash,  and  sodium  is  found  in  soda. 
Some  elements  have  been  named  after  the  country  in 
which  they  were  first  discovered.  Thus  we  have  gallium, 
discovered  in  France  ;  scandium,  discovered  in  Sweden  ; 
germanium,  discovered  in  Germany.  Tantalum  was  so 
called  on  account  of  the  long-continued  difficulties  ex- 
perienced in  isolating  it  when  it  was  discovered.  Colum- 
bium  received  its  name  from  the  fact  that  it  occurs  in 
the  mineral  columbite,  and  this  owes  its  name  to  the  fact 
that  it  was  first  found  in  the  United  States  of  America. 

Below  is  given  a  table  containing  the  names  of  all 
the  elementary  substances  now  known,  together  with 
their  symbols  and  atomic  weights.  The  names  of  those 
which  are  most  widely  distributed,  and  form  by  far  the 
largest  part  of  the  earth,  are  printed  in  SMALL  CAPITALS. 
The  names  of  those  which  are  rare  are  printed  in  italics. 


SYMBOLS  AND  ATOMIC  WEIGHTS  OF  THE  ELEMENTS.  21 


Element.  Symbol. 

ALUMINIUM Al 

Antimony Sb 

Argon A 

Arsenic As 

Barium Ba 

Bismuth Bi 

Boron B 

Bromine Br 

Cadmium Cd 

Ooesium Cs 

CALCIUM Ca 

CARBON C 

Cerium Ce 

CHLORINE Cl 

Chromium Cr 

Cobalt.. Co 

Columbium Cb 

Copper Cu 

Erbium E 

Fluorine F 

Gadolinium Gd 

Gallium Ga 

Germanium Ge 

Glucinum Gl 

Gold An 

Helium He 

HYDROGEN H 

'Indium In 

Iodine I 

Iridium Ir 

IRON Fe 

Lanthanum La 

Lead Pb 

Lithium Li 

MAGNESIUM Mg 

Manganese. . . , Mn 

Mercury Hg 


Atomic 
Weight. 
26.91 
119.52 

(?) 
74.44 

Element. 
Molybdenum  
Neodymium  
Nickel  
NITROGEN  .  .     .   . 

Symbol. 

....  Mo 
....  Nd 
....  Ni 

N 

136.39 

Osmium.. 

Os 

206.54 

OXYGEN  . 

o 

10.86 

Palladium.. 

Pd 

79.34 
111.10 

Phosphorus  
Platinum 

....  P 
Pt 

131.89 

POTASSIUM. 

K 

39.76 

Pi~aseodymium.  .  . 

Pr 

11.92 

JiJiodium 

Rh 

139.10 
35.18 

Rubidium  
Ruthenium  

....  Rb 
Ru 

51.74 

Samarium.. 

Sm 

58.49 

Scandium... 

Sc 

93.02 
63.12 

Selenium  
SILICON 

....  Se 
Si 

165.06 

Silver  

•  Ag 

18.91 

SODIUM  

.   ..  Na 

155.57 
69.38 

Strontium  
Sulphur  .  . 

....  Sr 
..  S 

71.93 

Tantalum 

Ta 

9.01 

Tellurium 

Te 

195.74 

Terbium.  . 

Tr 

(?) 
1.00 

Thallium  
Thorium  

....  Tl 
Th 

112.99 

Thulium  

Tm 

125.89 

Tin 

Sn 

191.66 

Titanium 

Ti 

55.60 

Tungsten 

W 

137  59 

Uranium 

u 

205.36 

Vanadium  

....  V 

697 

Ytterbium 

Yb 

24.10 

Yttrium.  

..  Y 

54.57 

Zinc  

.  ...  Zn 

198.49 

Zirconium..  . 

,.  Zr 

Atomic 
Weight. 

95.26 

139.70 

58.24 

13.93 

189.55 

15.88 

105.56 

30.79 

193.41 

38.82 

142.50 

,  102.23 

84.78 

100.91 

149.13 

43.78 

78.42 

28.18 

107.11 

22.88 

86.95 

31.83 

181.45 

126.52 

158.80 

202.61 

230.87 

169.40 

118.15 

47.79 

183.43 

237.77 

50.99 

171.88 

88.35 

64.91 

89.72 


The  symbols  of  the  elements  represent  not  only  the 
names  but  relative  quantities.  Thus  O  stands  for  15.88 
parts  by  weight  of  oxygen ;  N  for  13.93  parts  by  weight 
of  nitrogen ;  and  hydrogen  being  that  element  which 
enters  into  combination  in  the  smallest  relative  quantity, 
it  is  taken  as  the  basis  of  the  system,  or  H  stands 
for  1  part  by  weight  of  hydrogen.  What  the  symbol  O 
really  means  then  is  that  the  quantity  of  matter  repre- 
sented by  it  is  15.88  times  as  great  as  the  quantity  of 
matter  represented  by  the  symbol  H ;  and  the  quantity 
of  matter  represented  by  N  is  13.93  times  as  great  as  that 
represented  by  the  symbol  H ;  and  so  on  through  the 
list.  The  figures  do  not  represent  absolute  but  relative 
masses.  There  are  very  serious  difficulties  encountered 
in  determining  the  combining  weights  of  the  elements, 
and  in  regard  to  several  given  in  the  above  table  there  is 


22  INORGANIC  CHEMISTRY. 

considerable  doubt  as  to  the  accuracy.  Those  of  the  ele- 
ments with  which  we  most  frequently  have  to  deal  have, 
however,  been  determined  .with  great  care.  Work  in  this 
field  is  being  constantly  carried  on,  and  every  year  our 
knowledge  in  regard  to  the  combining  weights  becomes 
more  and  more  accurate. 

Symbols  of  Compounds.— As  the  elements  enter  into 
combination  in  the  proportion  of  their  respective  combin- 
ing weights  or  simple  multiples  of  these  weights,  it  is  an 
easy  matter  to  represent  the  composition  of  compounds 
by  means  of  the  symbols.  Thus  hydrogen  and  chlorine 
combine  in  the  proportion  of  their  combining  weights  to 
form  the  compound  hydrochloric  acid.  The  compound 
is  represented  by  the  symbol  HC1,  which  signifies  that 
the  compound  contains  hydrogen  and  chlorine  in  the 
proportion  of  35.18  parts  of  chlorine  to  1  part  of  hydro- 
gen. So  the  symbol  ZnO  means  a  chemical  compound 
consisting  of  64.91  parts  of  zinc  and  15.88  parts  of  oxygen ; 
HCN  means  a  compound  made  up  of  1  part  of  hydrogen, 
11.92  parts  of  carbon,  and  13.93  of  nitrogen.  Whenever 
the  symbols  of  the  elements  are  placed  side  by  side  with 
no  sign  between  them,  as  in  the  above  examples,  the  re- 
sulting symbol  means  that  the  elements  are  in  chemical 
combination.  But,  as  has  been  pointed  out,  elements 
may  combine  in  more  than  one  proportion.  In  one  of 
the  two  compounds  of  carbon  and  oxygen  the  elements 
are  combined  in  the  proportion  of  their  combining 
weights,  and  the  compound  is  represented  by  the  symbol 
CO ;  in  the  other  compound  the  elements  are  combined 
in  the  proportion  of  twice  the  combining  weight  of  oxy- 
gen to  the  combining  weight  of  carbon,  and  the  com- 
pound is  represented  by  the  symbol  CO2.  The  three 
compounds  of  iron  and  sulphur  to  which  reference  has 
already  been  made  are  represented  by  the  symbols  FeS, 
Fe2S3,  and  FeS2.  The  first  represents  a  compound  in 
which  the  elements  are  combined  in  the  proportion  of 
their  combining  weights,  or  55.60  parts  iron  to  31.83 
parts  sulphur ;  the  second  represents  a  compound  in 
which  the  elements  are  combined  in  the  proportion  of 
twice  the  combining  weight  of  iron  (2  X  55.60  =  111.20 


SYMBOLS  OF  COMPOUNDS- CHEMICAL  EQUATIONS.  23 

parts)  to  three  times  the  combining  weight  of  sulphur 
(3  X  31.83  =  95.49  parts) ;  and  the  third  represents  a  com- 
pound in  which  the  elements  are  combined  in  the  pro- 
portion of  the  combining  weight  (55.60  parts)  of  iron  to 
twice  the  combining  weight  (2  X  31.83  =  63.66  parts)  of 
sulphur.     The  four  compounds  of   potassium,  chlorine, 
and  oxygen  above  mentioned  are  represented  by  the  sym- 
bols KC1O,  KC1O2,   KC1O3,  and  KC1O4,  the  meaning  of 
which  will  be  clear  from  the  explanation  just  given.     By 
means  of  such  symbols  all  chemical  compounds  can  be 
represented,  and  they  represent  not  only  what  elements 
are  contained  in  the  compounds,  but  in  what  proportions 
the  elements  are  combined.     They  represent  facts  which 
have  been  determined  by  experiment.     Knowing  the  act- 
ual weight  of  one  constituent  of  any  compound  we  can 
calculate  by  the  aid  of  the  symbol  the  actual  weights  of 
the   other    constituents    and    of    the    compound   itself. 
Thus,  if  we  know  the  actual  weight  of  the  chlorine  con- 
tained in  a  quantity  of  potassium  chlorate,  KC1O3,  we 
can  calculate  how  much  potassium  and  how  much  oxygen 
are  contained  in  that  same  quantity,  and  also  what  the 
quantity  of  potassium  chlorate  is.     Suppose,  for  example, 
we  know  that  in  a  certain  quantity  of  potassium  chlorate 
there  is  contained  25  grams  of  chlorine,  and  it  is  desired 
to  know  how  much   potassium   and   how  much  oxygen 
there  are  in  this  quantity,  and  also  what  the  quantity  of 
potassium  chlorate  is.     We   know   that   the   compound 
KC1O8  is  made  up  of  38.82  parts  of  potassium,  35.18 
parts  of  chlorine,  and  3  times  15.88,  or  47.64,  parts  of 
oxygen,  the  whole  making  122.28  parts.     The   solution 
of  the  following  equations  in  proportion  will  give  the 
quantities  desired : 


35.18 
35.18 
35.18 


25         38.82  :  weight  of  potassium  ; 

25         47.64  :       "        "  oxygen ; 

25       121.64  :       "        "  potassium  chlorate. 

Chemical  Equations. — In  dealing  with  cases  of  chemi- 
cal action  it  is  desirable  to  express  by  means  of  the 
symbols  which  represent  the  elements  and  compounds 
what  takes  place.  In  general,  a  chemical  change  is  called 


24  INORGANIC  CHEMISTRY. 

a  chemical   reaction,  and   these   reactions  are  of  three 
kinds  : 

(1)  Two  or  more  elements  or  compounds  may  unite 
directly  to  form  one  product.  This  is  called  combination. 
The  following  examples  will  suffice.  When  mercury  is  kept 
boiling  in  the  air  for  a  time  it  becomes  covered  with  a  layer 
of  a  red  substance  which  is  a  compound  of  mercury  and 
oxygen  represented  by  the  symbol  HgO.  Magnesium 
burns  in  the  air  and  forms  the  compound  MgO.  Hydro- 
chloric acid,  HC1,  combines  directly  with  ammonia,  NH3, 
and  forms  the  compound  known  as  ammonium  chloride, 
NH4C1.  Water,  H2O,  combines  directly  with  lime  or  cal- 
cium oxide,  CaO,  to  form  slaked  lime  or  calcium  hydrox- 
ide, CaO2H2.  To  represent  these  facts,  the  symbols  of 
the  elements  or  compounds  which  act  upon  each  other 
are  written  with  a  plus  sign  between  them,  and  the  sign 
of  equality  is  written  before  the  symbol  of  the  product. 
The  chemical  equations  which  represent  the  above-men- 
tioned chemical  reactions  are  : 

Hg  +0       -HgO; 
Mg  +  O      =  MgO  ; 


H2O  +  CaO  =  CaO2H2. 

In  reading  the  equations  the  plus  sign  is  generally  ren- 
dered by  and,  and  the  sign  of  equality  by  give.  The  first 
equation  should  accordingly  be  read,  "  Mercury  and  oxy- 
gen give  mercuric  oxide  ;"  but  it  represents  besides  this 
fact  the  exact  relations  which  exist  between  the  quantities 
of  the  elements  and  the  compound  which  take  part  in  the 
reaction. 

(2)  The  second  kind  of  chemical  reaction  is  decomposi- 
tion or  the  opposite  of  combination.  Examples  are  fur- 
nished by  the  decomposition  of  mercuric  oxide  into  mer- 
cury and  oxygen  by  heat  ;  of  potassium  chlorate  into 
potassium  chloride  and  oxygen  by  heat  ;  of  water  into 
hydrogen  and  oxygen  by  the  electric  current  ;  and  of  cal- 
cium carbonate  or  limestone,  CaCO3,  into  lime  or  calcium 
oxide,  CaO,  and  carbon  dioxide,  CO2,  by  heat.  These  re- 
actions are  represented  by  the  following  equations  : 


CHEMICAL  EQUATIONS— THE  SCOPE  OF  CHEMISTRT.  25 

HgO    =Hg    +0; 
KC1O3  =  KC1  +  3O  ; 
H20      =  2H    +0; 
CaCO3  =  CaO  +  CO3. 

The  expressions  3O  and  2H  mean  respectively  three 
times  the  combining  weight  of  oxygen  and  twice  the  com- 
bining weight  of  hydrogen,  the  figure  being  generally 
used  in  this  way  when  the  element  is  not  in  combination. 
It  is,  however,  sometimes  written  the  same  as  if  the  ele- 
ment were  in  combination,  as  will  be  explained  later. 

(3)  The  third  kind  of  chemical  reaction  is  that  in 
which  two  or  more  compounds  give  rise  to  the  formation 
of  two  or  more  others ;  or  an  element  and  a  compound 
may  act  in  such  a  way  as  to  give  another  compound  and 
another  element.  This  is  called  double  decomposition 
or  metathesis.  The  following  cases  will  serve  as  exam- 
ples :  Sulphuric  acid,  H2SO4,  acts  upon  potassium  nitrate, 
KNO3,  or  saltpeter,  forming  potassium  sulphate,  K2SO4, 
and  nitric  acid,  HNO3;  nitric  acid,  HNO3,  acts  upon 
sodium  carbonate,  Na2CO3,  forming  sodium  nitrate, 
NaNO3,  carbon  dioxide,  CO2,  and  water,  H2O  ;  hydro- 
chloric acid,  HC1,  and  zinc,  Zn,  give  zinc  chloride,  ZnCl2, 
and  hydrogen.  These  facts  are  represented  as  below  : 

H2S04  +  2KN03  =  K2SO4    +  2HNO3; 
2HN03  +  Na2C03  =  2NaNO3  +  CO2  +  H2O  ; 
2HC1     +Zn          =ZnCl2      +2H. 

In  the  expressions  2KNO3,  2HNO3,  2NaNO3,  and  2HC1, 
the  large  figures  placed  before  the  symbol  of  the  com- 
pounds signify  that  the  quantities  of  the  compounds  repre- 
sented by  the  symbol  are  to  be  multiplied  by  the  figure. 
Thus,  HC1  stands  for  1  +  35.18  =  36.1.8  parts  of  hydro- 
chloric acid  ;  but  2HC1  stands  for  2(1  +  35.18)  =  72.36 
parts ;  3HC1  stands  for  3(1  +  35.18)  =  108.54  parts  ;  etc. 

The  Scope  of  Chemistry. — A  complete  study  of  chem- 
istry would  involve  the  study  of  the  action  of  all  the 
elements  upon  one  another  under  all  possible  circum- 
stances, and  a  study  of  the  action  of  all  compounds 
upon  one  another  and  upon  the  elements  under  all  cir- 


26  INORGANIC  CHEMISTRY. 

cumstances.  This  indicates  that  the  field  is  almost 
boundless ;  and  if  the  facts  were  not  related  among  one 
another,  if  every  time  a  reaction  is  studied  we  are  to  ex- 
pect something  entirely  different  from  all  other  reactions, 
the  task  would  be  practically  hopeless.  Fortunately  a 
great  many  general  facts  are  known,  and  reactions  which 
at  first  seem  to  have  no  connection  are  by  careful  study 
shown  to  be  related.  Thus  the  study  is  very  much  sim- 
plified and  made  interesting.  It  must  be  our  purpose  to 
study  the  facts  in  as  systematic  a  way  as  possible,  and  to 
be  constantly  on  the  alert  to  detect  relations.  The  habit 
of  comparing  a  new  reaction  with  others  already  studied 
should  be  cultivated.  In  this  way  light  will  come  out  of 
the  darkness,  and  the  subject  will  gradually  become  clear. 
While  the  simplest  way  to  begin  the  study  of  chemistry 
is  by  a  consideration  of  the  elements,  the  subject  is  com- 
plicated by  the  fact  that  we  cannot  readily  obtain  these 
elements  without  the  aid  of  substances  which  have  not 
been  studied,  and  of  processes  which  are  incomprehen- 
sible. There  are,  however,  two  elements  that  occur  in 
nature  in  enormous  quantities,  that  can  be  obtained  in 
the  iincombined  condition  quite  easily.  As  the  kinds 
of  action  which  they  exhibit  are  of  great  importance 
and  well  calculated  to  give  an  insight  into  the  nature  of 
chemical  action  in  general,  we  may  profitably  begin  our 
study  of  chemical  phenomena  by  a  study  of  these  two 
elements.  They  are  oxygen  and  hydrogen.  In  learning 
the  main  facts  in  regard  to  these  two  elements  we  shall 
learn  a  great  deal  that  will  be  of  importance  in  enabling 
us  to  understand  other  chemical  phenomena ;  we  shall 
begin  to  learn  how  to  study  things  chemically ;  and  we 
shall  thus  prepare  ourselves  for  a  systematic  study  of 
the  science  of  chemistry. 

Chemical  Action  accompanied  by  other  kinds  ol  Action. 
— Whenever  a  chemical  change  takes  place  it  is  accom- 
panied by  other  changes  and,  in  order  to  gain  a  com- 
plete knowledge  of  the  phenomenon,  these  other  changes 
must  be  studied.  Thus  when  sulphuric  acid  acts  upon 
zinc  the  chemical  change  is  represented  both  qualitatively 
and  quantitatively  by  the  equation 


CHEMICAL  REACTIONS.  21 

Zn  +  H2SO4  =  ZnSO4  +  2H. 

In  studying  the  reaction,  the  first  thing  to  do  is  to  learn 
the  nature  of  the  substances  formed,  and  the  relations 
between  the  substances  which  act  upon  each  other  and 
the  products.  This  may  be  called  the  purely  chemical 
study  of  the  reaction.  But  much  more  can  be  learned 
in  regard  to  it  by  careful  observation.  In  the  first  place, 
we  must  take  into  account  the  fact  that  a  solid  and  liquid 
here  react  to  form  a  solid  and  a  gas ;  and  the  ques- 
tion suggests  itself,  does  this  change  to  the  gaseous  con- 
dition exert  any  influence  on  the  reaction,  or  is  this  a  fact 
of  no  special  importance  ?  Again,  it  will  be  observed 
that  accompanying  the  chemical  change  there  is  a  marked 
rise  in  temperature,  and  we  naturally  inquire  whether  the 
quantity  of  heat  evolved  is*  definite  for  definite  quantities 
of  the  substances,  and,  if  so,  what  relation  exists  be- 
tween them.  There  are  still  other  changes  which  must  be 
taken  into  account  in  order  to  get  a  complete  knowledge 
of  a  chemical  reaction,  but,  as  yet,  the  study  of  the  other 
changes  has  not  been  taken  up  in  a  general  way,  and  our 
information  in  regard  to  them  is  comparatively  limited. 
Within  late  years  much  progress  has  been  made  in  the 
study  of  the  heat  changes  which  accompany  chemical 
changes.  It  has  been  found  that  every  chemical  change 
gives  rise  either  to  an  evolution  or  to  an  absorption  of 
heat,  and  that  for  definite  quantities  of  the  same  sub- 
stances under  the  same  circumstances  the  same  amount 
of  heat  is  evolved  or  absorbed.  The  special  study  of  the 
heat  changes  connected  with  chemical  changes  is  called 
thermochemistry.  A  consideration  of  the  facts  and  laws 
of  thermochemistry  is  of  assistance  in  dealing  with  chemi- 
cal reactions,  and  some  attention  will  be  paid  to  the  sub- 
ject in  this  book. 


CHAPTER  II. 

A  STUDY  OF  THE  ELEMENT  OXYGEN. 

Historical. — The  older  chemists  considered  air  to  be  a 
simple  substance,  but  the  experiments  of  Priestley  (1774) 
and  Scheele  (1775)  showed  that  the  air  contains  two  gases 
only  one  of  which  has  the  power  to  support  combustion ; 
and  they  succeeded  independently  of  each  other  in  show- 
ing that  oxygen  is  a  distinct  substance.  The  discovery 
of  oxygen  had  a  very  important  bearing  on  the  work  of 
Lavoisier  on  combustion,  and  it  was  he  who  gave  the 
name  oxygen  (or  oxygene)  to  the  gas,  for  the  reason  that 
he  supposed  it  to  be  the  essential  constituent  of  all  those 
chemical  substances  which  are  known  as  acids,  the  word 
being  derived  from  the  Greek  ogvcr,  acid,  and  yeveiv,  to 
produce.  While  this  is  generally  true,  it  has  since  been 
found  that  there  are  other  elements  which  have  the 
power  to  give  acid  properties  to  the  substances  into  the 
composition  of  which  they  enter,  and,  therefore,  the  name 
is  misleading. 

Occurrence. — Oxygen  is  the  most  widely  distributed 
and  most  abundant  element  of  the  earth.  It  forms,  as 
has  been  stated,  about  47  per  cent  of  the  solid  crust  of 
the  earth ;  eight-ninths  of  water ;  and  about  one-fifth 
of  the  air.  It  occurs  also  in  combination  with  carbon 
and  hydrogen,  or  with  carbon,  hydrogen,  and  nitrogen  in 
the  substances  which  go  to  make  up  the  structure  of  liv- 
ing things,  whether  vegetable  or  animal.  Besides  this  it 
forms  a  part  of  most  manufactured  chemical  products. 

Preparation. — Notwithstanding  the  abundant  supply  of 
oxygen  in  nature  it  is  not  a  simple  matter  to  get  it  in  the 
free  or  uncombined  state  from  most  substances  found  in 
nature.  As  it  forms  eight-ninths  of  water,  and  water 
consists  of  only  hydrogen  and  oxygen,  the  idea  suggests 
itself  at  once  that  it  may  be  made  by  the  decomposition 

(28) 


OXYGEN—  PREPARATION.  29 

of  water.  This  can  be  accomplished  without  serious 
difficulty  by  means  of  an  electric  current,  and  both  hy- 
drogen and  oxygen  obtained  in  this  way  ;  but  the  method 
is  expensive  and  more  complicated  than  others  which  are 
available,  and  therefore  it  is  not  commonly  used  for  the 
purpose.  In  the  air  the  two  gases  nitrogen  and  oxygen 
are  mixed  together  in  the  proportion  of  1  volume  of 
oxygen  to  4  volumes  of  nitrogen.  Here  then,  too,  as  in 
water,  we  have  an  enormous  supply,  but  it  is  difficult  to 
separate  the  oxygen  from  the  nitrogen  in  such  a  way  as 
to  leave  it  uncombined.  This  can,  however,  be  accom- 
plished, and  a  method  is  now  in  practical  use  on  the 
large  scale  for  the  purpose  of  preparing  oxygen  from  the 
air.  The  method  is  based  upon  the  fact  that  when 
barium  oxide,  BaO,  is  heated  in  a  current  of  air  it  takes 
up  oxygen  and  is  converted  into  barium  dioxide,  BaO2  ; 
and  when  the  pressure  upon  the  dioxide  is  sufficiently 
diminished,  it  is  decomposed  into  the  oxide,  BaO,  and 
oxygen,  as  represented  in  the  equation 


By  this  means  the  oxygen  can  be  extracted  from  the  air 
and  obtained  in  the  free  condition. 

Among  substances  which  occur  in  nature  and  which 
can  be  used  for  the  preparation  of  oxygen,  manganese 
dioxide  or  pyrolusite,  also  called  the  black  oxide  of  man- 
ganese, MnO2,  is  the  most  important.  It  gives  off  a  part 
of  its  oxygen  when  heated  to  a  comparatively  high  tem- 
perature. It  has  been  shown  that  the  decomposition  is 
represented  by  this  equation  : 


As  will  be  seen,  only  one-third  of  the  oxygen  contained 
in  the  dioxide  is  thus  obtained  in  the  free  state.  A  simi- 
lar method  is  that  used  by  Priestley  when  he  discovered 
oxygen.  It  consists  simply  in  heating  mercuric  oxide, 
HgO,  when  it  is  decomposed  as  represented  thus  : 

HgO  =  Hg  +  O. 

The  substance  potassium  chlorate,  KC1O3,  is  manufac- 
tured on  the  large  scale  for  a  variety  of  purposes  and  is, 


30  INORGANIC  CHEMISTRY. 

therefore,  easily  obtained.  It  gives  up  its  oxygen  when 
heated.  At  first  the  decomposition  represented  by  this 
equation  takes  place  : 

8KC103  =  5KC1O,  +  3KC1  +  4O. 

The  products  are  potassium  perchlorate,  KC1O4,  potas- 
sium chloride,  KC1,  and  oxygen,  one-sixth  of  the  total 
oxygen  being  given  off  in  the  first  stage.  This  part  of  the 
decomposition  takes  place  readily,  and  at  a  comparatively 
low  temperature.  If,  after  it  is  complete,  the  tempera- 
ture is  raised  considerably  higher,  more  gas  is  given  off 
and  the  change  represented  by  the  equation 

KC1O4  =  KC1  +  4O 

is  accomplished.  The  final  result  is,  therefore,  the  setting 
free  of  all  the  oxygen  contained  in  the  chlorate.  This- 
fact  is  represented  thus  : 

KC103  = 


The  best  method  for  use  in  the  laboratory  for  the 
preparation  of  oxygen  consists  in  heating  a  mixture  of 
equal  parts  of  coarsely  powdered  manganese  dioxide  and 
potassium  chlorate.  This  mixture  gives  up  oxygen  very 
readily  under  the  influence  of  heat.  Potassium  chlorate 
alone  requires  to  be  heated  to  a  temperature  of  over 
350°  C.  to  effect  its  decomposition,  but  when  mixed  with 
manganese  dioxide  the  decomposition  takes  place  at 
about  200°  C.  The  manganese  dioxide  does  not  lose  any 
of  its  oxygen  under  the  circumstances.  Other  substances, 
such  as  ferric  oxide,  copper  oxide,  etc.,  may  be  used  with 
similar  effect.  No  satisfactory  explanation  of  the  action 
of  these  substances  has  been  given.  Recent  experi- 
ments have  shown  that,  when  manganese  dioxide  is 
used,  oxygen,  chlorine,  and  potassium  permanganate, 
KMnO4,  are  first  formed.  The  permanganate  is  de- 
composed by  heat,  yielding  the  manganate,  K2MnO4, 
the  dioxide,  and  oxygen  ;  and,  finally,  the  manganate 
is  decomposed  by  chlorine,  yielding  potassium  chloride, 
the  dioxide,  and  oxygen. 

Physical  Properties.  —  Oxygen  is  a  colorless,  tasteless, 
inodorous  gas.     It  is  only  slightly  soluble  in  water,  100 


OXYGEN— CHEMICAL  PROPERTIES.  31 

volumes  of  water  at  0°  dissolving  4.1  volumes  of  oxygen. 
It  is  slightly  heavier  than  air,  its  specific  gravity  is 
1.10563,  and  1  liter  under  760  mm.  pressure  and  at  tem- 
perature 0°  weighs  1.429  grams,  while  a  liter  of  air 
weighs  1.2932  grams.  In  dealing  with  chemical  elements 
and  compounds  which  are  gaseous,  it  is  customary  to  use 
hydrogen  instead  of  the  air  as  the  standard  for  specific 
gravity.  While  the  specific  gravity  of  oxygen  in  terms 
of  air  is  1.10563,  in  terms  of  hydrogen  it  is  15.88,  or, 
in  other  words,  a  given  volume  of  oxygen  weighs  15.88 
as  much  as  the  same  volume  of  hydrogen  under  the 
same  conditions.  Under  a  pressure  of  50  atmospheres 
and  at  a  temperature  —  118°  it  is  condeDsed  to  a  liquid 
of  specific  gravity  0.978.  Liquid  oxygen  is  a  pale  steel- 
blue  transparent  and  very  mobile  liquid  which  boils 
at  — 181°.4  at  ordinary  pressure.  When  the  pressure  is 
reduced  or  removed,  evaporation  takes  place  so  rapidly 
that  a  part  of  the  oxygen  is  often  frozen  to  a  white  solid. 
Chemical  Properties. — At  ordinary  temperatures  oxy- 
gen does  not  act  readily  upon  most  other  things,  as  can 
be  clearly  shown  by  putting  a  variety  of  substances  in 
the  gas  without'heating  them.  If  they  are  left  for  a  con- 
siderable time  some  evidence  of  change  will  be  observed, 
but  generally  the  change  is  extremely  slow  unless  the 
temperature  is  raised.  At  higher  temperatures,  different 
for  different  substances,  it  combines  with  all  the  elements 
except  fluorine,  and  it  acts  readily  upon  a  large  number 
of  compounds.  Its  action  is  generally  accompanied  by 
an  evolution  of  heat  and  light,  and  the  process  under 
these  circumstances  is  called  combustion.  This  action 
may  be  illustrated  by  first  heating  and  then  introducing 
into  vessels  containing  oxygen,  sulphur,  charcoal,  iron  in 
the  shape  of  a  steel  watch-spring,  and  a  bit  of  phos- 
phorus. The  phenomena  observed  show  that  chemical 
action  takes  place,  but  they  do  not  show  what  is  formed. 
It  is  evident  that  in  each  case  light  and  heat  are 
evolved,  and  that  the  substances  introduced  into  the 
oxygen  are  changed  to  other  things.  In  the  case  of 
phosphorus  the  light  given  off  is  very  intense,  while  in 
that  of  carbon  and  that  of  sulphur  it  is  only  slight.  In 


32  INORGANIC  CHEMISTRY. 

the  vessel  in  which  the  burning  of  the  iron  takes 
place  a  reddish-brown  substance  is  deposited,  while  in 
that  in  which  the  phosphorus  is  burned  dense  white 
fumes  are  formed  and  at  first  the  product  is  partly  de- 
posited upon  the  walls  of  the  vessel  in  the  form  of  a 
white  powder  that  looks  like  snow.  After  standing  for 
some  time  over  water  it  disappears  and  the  water  evi- 
dently contains  something  in  solution.  A  thorough  study 
of  the  reactions  above  mentioned  has  shown  that  they 
consist  in  the  chemical  combination  of  oxygen  with  the 
substances  burned.  The  light  and  heat  are  results  of 
the  chemical  action.  The  reactions  are  represented  by 
the  following  equations : 

/t/        # 
With  sulphur,          8  +  2O  =  SOa ; 

"     carbon,  C  +  2O  =  CO, ; 

"     iron,  3Fe  +  4O  =  Fe3O4 ; 

"     phosphorus,  2P  +  5O  =  P2O5. 

The  products  are  sulphur  dioxide,  SO2,  a  colorless 
pungent  gas ;  carbon  dioxide,  CO2,  a  colorless  gas  ;  mag- 
netic oxide  of  iron,  Fe3O4,  a  reddish-brown  substance ; 
and  phosphorus  pentoxide,  PaO5,  a  white  solid  which  dis- 
solves in  water. 

Burning  in  the  Air  and  Burning  in  Oxygen. — One  can- 
not well  help  noticing  a  strong  resemblance  between  the 
burning  of  substances  in  the  air  and  in  oxygen  ;  and  the 
question  naturally  suggests  itself,  Are  these  two  acts 
the  same  in  character,  or  is  there  a  difference  between 
them  ?  To  answer  this  question,  we  must  burn  the  same 
substances  in  pure  oxygen  and  in  air,  and  determine 
whether  the  same  products  are  formed  in  the  two  cases, 
and  at  the  same  time  whether  anything  else  is  formed. 
If  we  should  make  this  comparison  in  any  case,  we 
should  find  that,  whether  a  substance  burns  in  the  air  or 
in  pure  oxygen,  the  same  product  is  formed,  and  nothing 
else.  It  is  therefore  certain  that  the  act  of  burning  in 
the  air  is  due  to  the  presence  of  oxygen.  But  substances 
do  not  burn  as  readily  in  the  air  as  in  oxygen,  and  some 
which  burn  in  oxygen  do  not  burn  in  the  air.  This  is 


PHLOGISTON  THEORY.  33 

due  to  the  fact  that  only  about  one-fifth  of  the  volume  of 
the  air  is  oxygen,  while  most  of  the  remaining  four-fifths 
consists  of  an  extremely  inactive  element,  nitrogen,  which 
takes  no  part  in  the  process  of  burning. 

Phlogiston  Theory.— Fire  in  its  various  forms  is  one  of 
the  longest  known  chemical  phenomena.  From  the  earli- 
est times  it  has  attracted  the  attention  of  men  and  has 
been  the  subject  of  speculative  and  experimental  study. 
It  was  one  of  the  elementary  substances  of  Aristotle,  as 
has  already  been  stated.  The  first  comprehensive  theory 
covering  all  cases  of  combustion  was  that  put  forward  by 
Stahl  and  known  as  the  pMogiston  theory.  According  to 
this,  every  combustible  substance  contains  something, 
called  phlogiston,  which  escapes  in  the  process  of  burn- 
ing. It  was  repeatedly  pointed  out  that  some  substances 
grow  heavier  by  burning,  or  rather  that  the  products  of 
combustion  weigh  more  than  the  substance  burned. 

This  can  be  shown  to  be  true  in  the  case  of  zinc  by 
placing  some  turnings  of  the  metal  on  one  pan  of  a  bal- 
ance and  determining  the  weight.  If  now  the  metal  is 
burned  it  will  change  almost  completely  to  a  white  pow- 
der, and  this  will  weigh  more  than  the  zinc. 

If  an  ordinary  candle  is  placed  on  one  pan  of  a  bal- 
ance and  a  glass  vessel  open  at  both  ends  and  filled  with 
large  pieces  of  sodium  hydroxide  or  caustic  soda  is  sus- 
pended directly  over  the  candle  in  such  way  that  the 
smoke  from  the  candle  must  pass  upwards  through  the 
tube,  and  a  similar  tube  is  suspended  over  the  other  pan 
of  the  balance,  and  equilibrium  is  then  established,  it 
will  be  found  that  the  pan  on  which  the  candle  is  placed 
will  gradually  grow  heavier  and  sink  as  the  candle  burns 
away.  The  explanation  of  this  is  simply  that  the  gases 
which  are  formed  by  the  burning  of  the  candle  are  re- 
tained by  the  caustic  soda,  and  they  weigh  more  than  the 
candle  from  which  they  were  formed  by  combustion. 

Facts  like  those  mentioned  were  known  when  the  phlo- 
giston theory  was  held,  but  they  do  not  appear  to  have 
made  a  very  strong  impression  upon  chemists.  Chemists 
were  at  all  events  not  able  to  give  a  satisfactory  explana- 
tion. This  remained  for  Lavoisier. 


34  INORGANIC  CHEMISTRY. 

Lavoisier's  Explanation  of  Combustion. — When  Lavoi- 
sier began  his  work  oxygen  was  unknown,  but  it  was 
discovered  soon  afterward  ;  and  this  discovery  was  of  the 
highest  importance  for  the  explanation  of  the  phenome- 
non of  combustion.  Lavoisier  showed  that  when  a  sub- 
stance is  burned  in  oxygen  or  in  the  air,  some  of  the  gas 
is  used  up.  He  then  weighed  the  substance  burned,  the 
oxygen  used  up,  and  the  product  formed,  and  found  that 
this  relation  holds  good  : 

Weight  of  Sub-      ,      Weight  of  Oxygen  Weight  of 

stance  burned  used  up  Product  formed. 

Having  established  this  relation  in  a  number  of  cases, 
it  followed  that  the  process  of  combustion  consists  in 
the  chemical  combination  of  oxygen  with  the  substance 
burned.  There  was  no  longer  room  for  the  hypotheti- 
cal phlogiston,  and  since  that  time  it  has  not  occupied 
a  place  in  the  thoughts  of  most  chemists. 

Combustion. — By  the  term  combustion  in  its  broadest 
sense  is  meant  any  chemical  act  which  is  accompanied 
by  an  evolution  of  light  and  heat.  Ordinarily,  however, 
it  is  restricted  to  the  union  of  substances  with  oxygen,  as 
this  union  takes  place  in  the  air,  with  evolution  of  light 
and  heat.  Substances  which  have  the  power  to  unite 
with  oxygen  are  said  to  be  combustible,  and  substances 
which  have  not  this  power  are  said  to  be  incombustible. 
Most  elements  combine  with  oxygen  under  proper  con- 
ditions, and  are  therefore  combustible.  Most  compounds 
formed  by  the  union  of  oxygen  with  combustible  sub- 
stances are  incombustible.  For  example,  the  sulphur 
dioxide,  carbon  dioxide,  magnetic  oxide  of  iron,  and  phos- 
phorus pentoxide,  formed  when  sulphur,  carbon,  iron,  and 
phosphorus  are  burned  in  oxygen,  are  incombustible. 
They  contain  oxygen  and  they  cannot  combine  directly 
with  more. 

Kindling  Temperature. — We  have  seen  that  substances 
do  not  usually  combine  with  oxygen  at  ordinary  tempera- 
tures, but  that  in  order  to  effect  the  union  the  tempera- 
ture must  be  raised.  If  this  were  not  the  case,  it  is  plain 
that  every  combustible  substance  in  nature  would  burn  up, 


KINDLING  TEMPERATURE— SLOW  OXIDATION.       35 

for  the  air  probably  supplies  a  sufficient  quantity  of  oxygen 
for  this  purpose.  Some  substances  need  to  be  heated  to  a 
high  temperature  before  they  will  combine  with  oxygen  ; 
others  require  to  be  heated  but  little.  If  we  were  to  sub- 
ject pieces  of  phosphorus,  of  sulphur,  and  of  carbon  to 
the  same  gradual  rise  in  temperature,  we  should  find 
that  the  phosphorus  takes  fire  very  easily,  only  a  slight 
elevation  of  temperature  being  necessary  ;  next  in  order 
would  come  the  sulphur ;  and  last  the  carbon.  If  we 
were  to  repeat  these  experiments  a  number  of  times,  we 
should  find  that  the  phosphorus  always  takes  fire  at  the 
same  temperature,  and  a  similar  result  would  be  reached 
in  the  case  of  sulphur  and  carbon.  Every  combustible 
substance  has  its  kindling  temperature ;  that  is,  the  tem- 
perature at  which  it  will  combine  with  oxygen.  Below 
this  temperature  it  will  not  combine  with  oxygen.  If  a 
piece  of  wood  should  be  heated  to  its  kindling  tempera- 
ture all  at  once,  it  would  burn  up  as  rapidly  as  it  could 
secure  the  necessary  oxygen ;  but  the  burning  does  not 
usually  take  place  rapidly,  for  the  reason  that  only  a 
small  part  of  it  is  at  any  one  time  heated  to  the  kind- 
ling temperature.  Watch  a  stick  of  wood  burning,  and 
see  how,  as  we  say,  "  the  fire  creeps  "  slowly  along  it. 
The  reason  of  the  slow  advance  is  simply  this :  Only 
those  parts  of  the  stick  which  are  nearest  the  burning 
part  become  heated  to  the  kindling  temperature.  They 
take  fire  and  heat  the  parts  nearest  them,  and  so  on 
gradually  throughout  the  length  of  the  stick. 

Slow  Oxidation.— Substances  may  combine  slowly  with 
oxygen  without  evolution  of  light.  Thus,  if  a  piece  of 
iron  is  allowed  to  lie  in  moist  air,  it  becomes  covered 
with  rust.  The  rust  is  similar  to  the  substance  formed 
when  iron  is  burned  in  oxygen.  Both  are  formed  by  the 
union  of  iron  and  oxygen.  Magnesium  burns  in  the  air 
and  forms  a  white  compound  containing  oxygen.  It 
burns  with  increased  brilliancy  in  oxygen,  forming  the 
same  compound.  If  left  in  moist  air  for  some  days  or 
weeks,  it  becomes  covered  with  a  layer  of  the  same  white 
substance.  If  this  is  scraped  off,  and  the  magnesium 
again  allowed  to  lie,  it  will  again  become  covered 


36  INORGANIC  CHEMISTRY. 

with  a  layer  of  the  compound  with  oxygen,  and  this  may 
be  continued  until  the  magnesium  has  been  completely 
converted  into  the  same  substance  that  is  formed  when 
it  burns  in  oxygen  or  in  the  air.  Many  other  cases  of 
slow  oxidation  might  be  described,  some  of  which,  such 
as  the  decay  of  wood,  are  constantly  taking  place.  The 
most  important  illustration  of  slow  oxidation  is  that 
which  takes  place  in  our  bodies,  for,  as  we  shall  see,  the 
food  of  which  we  partake  undergoes  a  great  many 
changes,  some  of  the  substances  uniting  with  oxygen, 
and  thus  keeping  up  the  temperature  of  our  bodies. 
This,  however,  is  done  without  evolution  of  light.  We 
take  large  quantities  of  oxygen  into  our  lungs  in  the  act 
of  breathing.  This  acts  upon  various  substances  pre- 
sented to  it,  oxidizing  them  to  other  forms  which  can 
easily  be  got  rid  of.  More  will  be  said  in  regard  to  the 
breathing  process  of  animals  and  plants  when  the  sub- 
ject of  carbon  and  its  compounds  is  taken  up. 

Heat  of  Combustion.— What  is  the  chief  difference  be- 
tween combustion,  as  we  ordinarily  understand  it,  and 
slow  oxidation?  As  far  as  can  be  judged  by  a  cursory 
examination,  it  is  that  in  the  former  there  is  an  evolution 
of  light  and  much  heat,  while  in  the  latter  there  is  ap- 
parently but  little  heat  evolved  and  no  light.  Remem- 
bering that  the  reason  why  a  body  gives  light  is  that  it 
is  heated  to  a  sufficiently  high  temperature,  the  problem 
resolves  itself  into  a  question  of  heat.  What  difference, 
if  any,  is  there  between  the  quantity  of  heat  given  off 
when  a  substance  burns,  and  when  it  undergoes  slow 
oxidation  without  evolution  of  light  ?  Experiment  has 
shown  that  there  is  no  difference.  In  one  case  all  the  heat 
is  given  off  in  a  short  space  of  time,  and  therefore  the 
temperature  of  the  substance  becomes  high  and  it  emits 
light.  In  the  other  the  heat  is  given  off  slowly  and  con- 
tinues for  a  much  longer  time,  and  therefore  the  tem- 
perature of  the  substance  does  not  get  high,  as  surround- 
ing substances  conduct  off  the  heat  nearly  as  rapidly  as 
it  is  evolved.  If,  however,  we  were  to  measure  the 
quantity  of  the  heat,  we  should  find  it  to  be  the  same  in 
both  cases. 


HEAT  OF  COMBUSTION.  37 

We  can  measure  the  heat  given  off  or  absorbed  in  a 
chemical  reaction  by  allowing  the  reaction  to  take  place 
in  a  vessel  called  a  calorimeter,  so  constructed  as  to 
prevent  loss  of  heat,  and  containing  a  known  weight  of 
water.  The  temperature  of  the  water  is  noted  at  the 
beginning  of  the  operation  and  at  the  end.  A  quantity 
of  heat  is  generally  stated  by  giving  the  number  of  grams 
of  water  which  it  will  raise  one  degree  (Centigrade)  in  tem- 
perature. The  quantity  of  heat  necessary  to  raise  a  gram 
of  water  one  degree  in  temperature  is  the  unit  used  in  heat 
measurements.  It  is  called  the  calorie.  If  we  say  that 
the  quantity  of  heat  evolved  in  any  reaction  is  250  cal- 
ories (written  generally  250  cal.),  this  means  simply  a 
quantity  of  heat  capable  of  raising  the  temperature  of 
250  grams  of  water  one  degree  or  of  one  gram  of  water 
250  degrees  in  temperature.  Sometimes  it  is  convenient 
to  use  a  larger  unit.  The  quantity  of  heat  required  to 
raise  the  temperature  of  one  kilogram  of  water  one  de- 
gree serves  the  purpose.  This  is  the  large  calorie.  To 
distinguish  it  from  the  smaller  one  it  is  written  with  a 
capital.  Thus,  250  Cal.  means  250  large  calories.  The 
large  calorie  is  obviously  1000  times  greater  than  the 
small  calorie. 

To  repeat,  then :  by  the  heat  of  combustion  of  a  sub- 
stance is  meant  simply  the  quantity  of  heat  given  off 
when  a  certain  weight  of  the  substance  combines  with 
oxygen.  In  order  to  avoid  confusion  it  is  necessary  to 
have  an  agreement  in  regard  to  the  weight  of  substance 
which  shall  be  used  as  the  standard.  This  may  be  a 
gram  or  any  other  weight,  but  for  the  purposes  of  chem- 
istry it  is  most  convenient  to  take  weights  in  proportion 
to  the  combining  or  atomic  weights.  Thus,  by  the  heat 
of  combustion  of  carbon  is  meant  the  quantity  of  heat 
evolved  by  the  combination  of  11.92  grams  of  carbon 
with  2  X  15.88  =  31.76  grams  of  oxygen.  By  the  heat 
of  combustion  of  sulphur  is  meant  the  quantity  of  heat 
evolved  in  the  combination  of  31.83  grams  sulphur  with 
2  X  15.88  =  31.76  grams  oxygen,  etc. 

Not  only  is  the  heat  of  combustion  the  same  whether 


38  INORGANIC  CHEMISTRY. 

the  union  with  oxygen  takes  place  slowly  or  rapidly, 
but  the  heat  evolved  in  any  given  chemical  reaction  is 
always  the  same,  and  chemical  action  is  always  accom- 
panied by  an  evolution  or  absorption  of  heat. 

Heat  of  Decomposition.— Just  as  it  is  true  that  a  definite 
quantity  of  heat  is  evolved  when  two  or  more  elements 
combine  chemically,  so  also  it  is  true  that  in  order  to 
overcome  the  force  which  holds  these  elements  together 
the  same  quantity  of  heat  is  absorbed.  Thus,  the  heat 
of  formation  of  mercuric  oxide,  HgO,  is  30,660  cal.;  or, 
in  other  words,  when  198.49  grams  of  metallic  mercury 
and  15.88  grams  of  oxygen  combine,  30,660  calories  of 
heat  are  evolved.  Now,  we  have  seen  that  when  heat  is 
applied  to  the  compound  it  is  decomposed  into  its 
elements.  To  effect  this  decomposition,  as  much  heat 
is  absorbed  as  was  evolved  in  the  formation  of  the  com- 
pound. 

Chemical  Energy  and  Chemical  Work.— Any  substance 
which  has  the  power  to  unite  with  others  can  do  chemical 
work :  it  possesses  chemical  energy.  Thus,  all  combustible 
substances  can  do  work.  In  uniting  with  oxygen  heat 
is  evolved,  and  this  can  be  transformed  into  motion.  In 
the  case  of  the  steam-engine,  the  cause  of  the  motion  is 
the  burning  of  the  fuel,  which  is  a  chemical  act.  We 
thus  see  that  the  source  of  the  power  of  the  steam-engine 
is  chemical  energy.  Substances,  on  the  other  hand, 
which  have  no  power  to  combine  with  others  have  no 
power  to  do  chemical  work,  or  they  have  no  chemical 
energy.  So  far  as  power  to  combine  with  oxygen  is  con- 
cerned, water  is  a  substance  of  this  kind,  as  is  also  car- 
bon dioxide,  the  gas  formed  when  carbon  is  burned  in 
oxygen.  In  order  that  they  may  do  work  by  combining 
with  oxygen,  they  must  first  be  decomposed,  and  their 
constituents  put  together  in  some  form  in  which  they 
have  the  power  of  combination.  This  decomposition 
of  carbon  dioxide  and  water  is  taking  place  constantly 
on  the  earth.  All  plant-life  is  dependent  on  it.  The 
products  of  the  action,  i.e.,  the  different  kinds  of  wood, 
have  chemical  energy, — they  can  do  chemical  work. 
This  power  to  do  work  has  been  acquired  from  the  heat 


OXIDES.  39 

of  the  sun,  which  is  the  main  force  used  in  decomposing 
the  carbon  dioxide  and  water.  We  have  thus  a  trans- 
formation of  the  sun's  heat  into  chemical  energy,  which 
is  stored  up  in  the  combustible  woods.  The  quantity 
of  heat  which  is  given  off  in  burning  wood  is  believed  to 
be  exactly  equal  to  the  quantity  of  heat  used  up  in  its 
formation* 

Oxides.— The  compounds  of  oxygen  with  other  elements 
are  called  oxides.  To  distinguish  between  different  ox- 
ides, the  name  of  the  element  with  which  the  oxygen  is 
in  combination  is  prefixed.  Thus,  the  compound  of  zinc 
and  oxygen  is  called  zinc  oxide;  that  of  calcium  and 
oxygen,  calcium  oxide;  that  of  silver  and  oxygen,  silver 
oxide;  etc.  When  an  element  forms  more  than  one  com- 
pound with  oxygen,  suffixes  are  used  to  distinguish 
between  them.  Thus  in  the  case  of  copper  there  are 
two  oxides  which  have  the  composition  represented  by 
the  symbols  Cu2O  and  CuO.  The  former  is  known  as 
cuprous  oxide  and  the  latter  as  cupric  oxide.  That  oxide 
which  contains  the  smaller  quantity  of  oxygen  in  com- 
bination with  a  given  quantity  of  the  other  element  is 
designated  by  the  suffix  ous ;  that  which  contains  the 
larger  proportion  of  oxygen  is  designated  by  the  suffix 
ic.  In  other  cases  the  number  of  combining  weights  of 
oxygen  contained  in  the  compound  is  indicated  by  the 
name.  Thus,  manganese  dioxide  is  MnO2;  sulphur  tri- 
oxide  is  SO3;  etc. 


CHAPTER  III. 

A  STUDY  OF  THE  ELEMENT  HYDROGEN. 

Historical.— Hydrogen  was  discovered  as  a  distinct 
substance  by  Cavendish  in  1766,  although  it  had  been 
observed  as  an  inflammable  gas  before  that  time. 

Occurrence.— It  occurs  to  some  extent  in  the  free  con- 
dition, and  issues  from  the  earth  in  small  quantity  in 
some  localities.  It  is,  for  example,  a  constituent  of 
the  gases  which  escape  from  the  petroleum  wells  in 
Pennsylvania.  It  has  also  been  shown  to  occur  in  enor- 
mous quantities  in  the  atmosphere  of  the  sun.  On  the 
earth  it  occurs  chiefly,  however,  in  combination  in  water, 
of  which  it  forms  11.11  per  cent.  It  occurs  also  in  most 
substances  of  animal  and  vegetable  origin,  such  as  the 
various  kinds  of  wood  and  fruits,  and  the  tissues  of  all 
animals.  In  these  products  of  life  it  is  contained  in 
combination  with  carbon  and  oxygen  or  with  carbon, 
oxygen,  and  nitrogen. 

Preparation.— The  simplest  way,  theoretically,  to  pre- 
pare hydrogen  is  by  the  decomposition  of  water  by  the 
electric  current.  It  has  already  been  stated  that  when 
an  electric  current  is  passed  through  water  the  two  gases 
oxygen  and  hydrogen  are  liberated.  But  this  method 
is  less  convenient  and  more  expensive  than  other 
methods  which  are  available,  and  it  is  therefore  used 
only  under  special  circumstances.  It  is  particularly 
well  adapted  to  the  preparation  of  small  quantities  of 
pure  hydrogen. 

Some  elements  when  brought  in  contact  with  water  at 
the  ordinary  temperature  decompose  it  and  set  hydrogen 
free.  The  two  most  easily  obtained  elements  which  act 
in  this  way  are  sodium  and  potassium.  If  a  small  piece 
of  potassium  is  thrown  upon  water,  a  flame  is  observed 
at  once.  If  sodium  is  used,  it  is  seen  to  form  a  small 

(40) 


HYDROGEN-— PREPARATION.  41 

ball  which  moves  about  on  the  surface  of  the  water  with 
a  hissing  sound,  but  under  ordinary  circumstances  no 
flame  is  observed.  By  applying  a  flame  to  the  ball 
something  takes  fire  and  burns.  By  filling  a  good-sized 
test-tube  with  water  and  inverting  it  in  a  larger  vessel 
and  bringing  a  small  piece  of  sodium  wrapped  in  a  piece 
of  filter-paper  below  the  mouth  of  the  tube,  the  sodium 
will  rise  to  the  top  of  the  tube  when  released,  and  it 
will  then  be  seen  that  a  gas  is  evolved  which  gradually 
depresses  the  water  in  the  tube.  A  similar  experi- 
ment with  potassium  gives  a  similar  result.  The  gas 
given  off  in  each  case  is  hydrogen.  By  evaporating  off 
the  water  left  in  the  vessel  there  will  be  found  in  each 
case  a  white  substance  of  marked  chemical  properties. 
That  formed  with  the  potassium  is  known  as  potassium 
hydroxide,  or  caustic  potash,  and  has  the  composition 
represented  by  the  symbol  KOH ;  that  formed  with  the 
sodium  is  known  as  sodium  hydroxide,  or  caustic  soda, 
and  is  represented  by  the  symbol  NaOH.  The  reactions 
which  take  place  between  potassium  and  sodium  and 
water  are  represented  by  the  equations 

E  +  H2O  =  KOH   +  H ;  and 
Na  +  H20  =  NaOH  +  H. 

Half  the  hydrogen  of  the  water  which  is  decomposed  is 
replaced  by  the  potassium  or  the  sodium,  as  the  case 
may  be.  These  reactions  are  partly  described  by  say- 
ing that  the  potassium  or  sodium  is  substituted  for  half 
the  hydrogen  in  the  water,  and  the  act  is  called  sub- 
stitution. This  is  a  very  common  kind  of  chemical  ac- 
tion, and  we  shall  constantly  meet  with  it  in  the  course  of 
our  study.  The  reaction  is  one  of  double  decomposition 
or  metathesis,  two  substances  acting  upon  each  other  to 
form  two  others.  The  cause  of  such  a  reaction  is  to  be 
sought  for  in  the  different  degrees  of  attraction  exerted 
by  the  elements  upon  one  another.  In  general  terms,  if 
two  compounds  AB  and  CD  are  brought  together,  and 
the  element  A  has  for  C  a  stronger  attraction  than  A  for 
B,  and  B  has  for  D  a  stronger  attraction  than  it  has  for 


42  INORGANIC  CHEMISTRY. 

A,  then  reaction  will  take  place,  to  some  extent  at  least, 
according  to  the  equation 

AB  +  CD  =  AC+BD. 


The  action  may  be  modified  by  a  number  of  circum- 
stances which  will  be  treated  of  in  due  time.  It  is  evi- 
dent, however,  now  that  a  very  important  problem  for 
the  chemist  to  solve  is  the  determination  of  the  attrac- 
tion which  the  elements  exert  upon  one  another.  It  is 
extremely  difficult  to  make  these  determinations,  but 
something  can  be  learned  in  regard  to  them  by  a  study 
of  the  changes  in  temperature  which  accompany  them. 
The  attraction  is  not  proportional  to  the  heat  evolved,  for 
reasons  which  will  be  pointed  out  later,  but  there  is  some 
relation  between  them. 

Some  substances  which  decompose  water  slowly  at  the 
ordinary  temperature  do  so  readily  at  a  higher  tempera- 
ture. This  is  true,  for  example,  of  iron.  At  ordinary 
temperatures  it  decomposes  water,  as  is  seen  in  the  for- 
mation of  a  coating  upon  it  when  left  in  contact  with 
water.  At  higher  temperatures  when  the  iron  is  red- 
Lot  it  decomposes  water  very  readily,  and  hydrogen  may 
be  made  in  quantity  by  this  means.  In  the  laboratory 
the  iron  may  be  heated  in  a  gun-barrel  or  in  a  porcelain 
tube.  When  steam  is  passed  over  it  the  decomposition 
represented  in  this  equation  takes  place  : 

3Fe  +  4H2O  =  Fe3O4  +  8H. 

The  iron  combines  with  the  oxygen  and  liberates  the 
hydrogen. 

Carbon,  in  the  form  of  charcoal  or  coal,  may  be  used 
in  a  similar  way  to  effect  the  decomposition  of  water 
and  the  liberation  of  hydrogen.  At  a  high  heat  the  re- 
action takes  place  mainly  as  represented  thus  : 

C  +  H20  =  CO  +  2H. 

A  mixture  of  two  gases,  carbon  monoxide  and  hydrogen, 
is  thus  formed.  This  mixture  is  the  essential  part  of  the 
gas  which  has  of  late  years  come  into  such  extensive 


HYDROGEN—  PREPARATION.  43 

use  under  the  name  "  water-gas."  This  is  formed  by 
passing  steam  over  highly  heated  anthracite  coal. 

By  far  the  most  convenient  method  for  making  hydro- 
gen consists  in  treating  a  metal  with  an  acid.  Among 
the  metals  best  adapted  to  the  purpose  are  zinc  and 
iron,  and  indeed  zinc  is  almost  exclusively  used.  As 
will  be  seen  later,  acids  are  substances  that  contain 
hydrogen,  and  are  characterized  by  the  property  that 
they  give  up  this  hydrogen  very  easily  and  take  up 
other  elements  in  the  place  of  it.  Among  the  common 
acids  found  in  every  laboratory  are  hydrochloric  acid, 
sulphuric  acid,  and  nitric  acid.  The  chemistry  of  these 
compounds  will  be  treated  of  in  due  time  ;  but,  as  we 
shall  be  obliged  to  use  them  before  they  are  taken  up 
systematically,  a  few  words  in  regard  to  them  are  desir- 
able in  this  place. 

Hydrochloric  acid  is  a  compound  of  hydrogen  and 
chlorine.  It  is  a  gas  which  dissolves  easilv  in  water. 

t/ 

It  is  this  solution  which  is  used  in  the  laboratory,  and 
which  is  manufactured  in  enormous  quantities  in  connec- 
tion with  the  manufacture  of  soda  or  sodium  carbonate. 
Its  chemical  symbol  is  HC1.  In  commerce  it  is  not 
uncommonly  called  "  muriatic  acid." 

Sidphuric  acid  is  a  compound  of  sulphur,  oxygen,  and 
hydrogen  in  the  proportions  represented  by  the  formula 
H2SO4.  It  is  an  oily  liquid  and  is  frequently  called 
"  oil  of  vitriol."  It  is  manufactured  in  very  large  quan- 
tities, as  it  plays  an  important  part  in  many  of  the  most 
important  chemical  industries. 

Nitric  acid  is  a  compound  containing  nitrogen,  oxygen, 
and  hydrogen  in  the  proportions  represented  by  the 
formula  HNO3.  It  is  a  colorless  liquid,  though,  as  we 
get  it,  it  is  commonly  colored  straw-yellow. 

When  a  metal,  such  as  zinc,  is  brought  in  contact  with 
hydrochloric  or  sulphuric  acid,  an  evolution  of  hydro- 
gen takes  place  at  once.  The  reactions  are  as  repre- 
sented in  these  equations  : 


Zn  +  2HCl    =  ZnCl2  +2H; 
Zn  +  H2SO4  =  ZnSO4  +  2H. 


44  INORGANIC  CHEMISTRY. 

Each  combining  weight  of  zinc  liberates  and  replaces 
two  combining  weights  of  hydrogen. 

The  action  between  iron  and  these  two  acids  is  of  the 
same  character  : 


Fe  +  2HCl    =  FeCl2   +2H; 
Fe  +  H2SO4  =  FeSO4  +  2H. 

The  hydrogen  obtained  from  acids  by  the  action  of 
metals  is  not  pure,  but  it  can  be  purified  by  treatment 
with  appropriate  substances.  That  obtained  by  the  de- 
composition of  water  by  the  electric  current  is  pure. 

Physical  Properties.  —  Hydrogen  is  a  colorless,  inodor- 
ous, tasteless  gas.  That  made  by  the  action  of  acids  on 
zinc  or  iron  has  a  somewhat  disagreeable  odor  which  is 
due  to  the  presence  of  other  gases  in  small  quantity.  It 
is  not  poisonous,  and  may  therefore  be  inhaled  with  im- 
punity. We  could  not,  however,  live  in  an  atmosphere 
of  hydrogen,  as  we  need  oxygen.  It  is  the  lightest 
known  substance.  Its  specific  gravity  in  terms  of  the 
air  standard  is  0.06926.  A  litre  under.  760  mm.  pressure 
and  at  0°  weighs  0.089873  gram.  Under  Oxygen  it  was 
stated  that  in  chemistry  hydrogen  is  commonly  taken  as 
the  standard  of  specific  gravity,  and  that,  hydrogen  being 
unity,  the  specific  gravity  of  oxygen  is  15'.  88.  The  gas  is 
only  slightly  soluble  in  water.  100  volumes  of  water 
take  up  1.93  volumes  of  hydrogen.  The  fact  that  hy- 
drogen is  lighter  than  the  air  is  shown  by  opening  a 
vessel  which  contains  it  and  turning  the  mouth  of  the 
vessel  upward.  The  gas  escapes  at  once,  and  in  a  very 
short  time  no  evidence  of  its  presence  can  be  obtained. 
Light  vessels  as,  for  example,  soap-bubbles  or  collodion- 
balloons  filled  with  the  gas  rise  in  the  air,  and  it  is  used 
for  the  purpose  of  filling  large  balloons. 

Hydrogen  has  been  converted  into  the  liquid  form 
under  a  pressure  of  20  atmospheres  at  a  temperature  of 
—  234°.  5.  Its  boiling-point  under  ordinary  atmospheric 
pressure  is  —  243°.5. 

Hydrogen  passes  readily  through  porous  substances, 
or  it  diffuses  rapidly.  This  can  easily  be  demonstrated 
in  the  case  of  porous  earthenware  and  paper.  It 


HYDROGEN-CHEMICAL  PROPERTIES.  45 

also  passes  readily  through,  some  metals,  as  iron  and 
platinum,  when  heated  to  redness.  There  is  a  direct  re- 
lation between  the  specific  gravity  of  gases  and  the  rate 
at  which  they  diffuse.  The  lower  the  specific  gravity 
the  more  rapid  the  diffusion.  The  law  governing  these 
phenomena  is  : 

The  rate  of  diffusion  of  gases  is  approximately  inversely 
proportional  to  the  square  roots  of  their  specific  gravities. 

The  specific  gravity  of  hydrogen  being  1  and  that  of 
oxygen  nearly  16  (15.88),  the  rate  of  diffusion  of  oxy- 
gen is  approximately  ^  that  of  hydrogen.  If  hydrogen 
is  on  one  side  of  a  porous  wall,  and  oxygen  on  the 
other,  the  hydrogen  will  pass  through  the  wall  so 
much  more  rapidly  than  the  oxygen  that  there  will  be 
an  accumulation  of  hydrogen  on  one  side  of  the  wall, 
and  if  the  vessel  were  closed  there  would  be  in- 
creased pressure  on  that  side.  The  ready  passage  of 
gases  through  porous  walls  is  a  matter  of  great  impor- 
tance in  connection  with  the  ventilation  of  dwellings. 
Most  of  the  materials  used  in  building  are  porous  and 
permit  the  passage  of  gases  through  them  in  both  direc- 
tions, and  change  of  air  is  secured  in  this  way  to  some 
•extent. 

Chemical  Properties. — Under  ordinary  circumstances, 
hydrogen  is  not  a  particularly  active  element.  It  does 
not  unite  with  oxygen  gas  at  ordinary  temperatures, 
but,  like  other  combustible  substances,  it  must  be 
heated  up  to  the  kindling  temperature  before  it  will 
l)urn.  If  a  lighted  match  is  applied  to  it,  it  takes 
fire  at  once.  The  flame  is  colorless  or  slightly  blue. 
Generally  the  flame  is  somewhat  colored  in  consequence 
of  the  presence  of  foreign  substances ;  but  that  it  is 
colorless  when  the  gas  is  burned  alone  can  be  shown  by 
burning  it  as  it  issues  from  a  platinum  tube  which  is 
itself  not  chemically  acted  upon  by  the  heat.  Although 
the  flame  is  not  luminous  it  is  intensely  hot,  as  can  be 
.seen  by  inserting  into  it  a  coil  of  platinum  wire,  which 
will  at  once  become  red-hot  and  emit  light  accordingly. 

The  burning  of  hydrogen  in  the  air,  like  the  burning 
•of  other  combustible  substances  in  the  air,  consists  in  a 


46  INORGANIC  CHEMISTRY. 

union  of  the  gas  with  oxygen.  This  has  been  shown  to 
be  true  by  most  elaborate  experiments  on  the  combus- 
tion of  hydrogen  in  oxygen  and  in  the  air.  On  the  other 
hand,  substances  which  burn  in  the  air  are  extinguished 
when  put  in  a  vessel  containing  hydrogen.  This  is- 
equivalent  to  saying  that  a  substance  which  is  uniting 
with  oxygen  does  not  continue  to  unite  with  oxygen 
when  put  in  an  atmosphere  of  hydrogen,  and  does  not 
combine  with  hydrogen.  The  fact  is  expressed  by  say- 
ing that  hydrogen  does  not  support  combustion.  Thi& 
can  be  shown  by  holding  a  vessel  filled  with  hydrogen, 
with  the  mouth  downward,  and  inserting  into  it  a  lighted 
taper  supported  on  a  wire.  The  gas  takes  fire  at  the 
mouth  of  the  vessel,  but  the  taper  is  extinguished. 

Ordinarily  we  say  that  hydrogen  burns  in  oxygen,, 
but,  as  the  act  consists  in  the  union  of  the  two  gases,  it 
would  seem  probable  that  oxygen  will  burn  in  an  at- 
mosphere of  hydrogen.  This  can  be  shown  to  be  true 
by  a  proper  arrangement  of  apparatus.  If  we  were 
surrounded  by  an  atmosphere  of  hydrogen  we  should 
probably  speak  of  oxygen  as  a  combustible  gas  in  the 
same  way  that  we  now  speak  of  hydrogen  as  a  combus- 
tible gas. 

It  can  easily  be  shown  that,  when  hydrogen  is  burned 
either  in  oxygen  or  air,  water  is  formed.  The  simplest 
way  to  show  this  is  by  holding  a  glass  plate  or  some 
other  incombustible  object  a  short  distance  above  a 
flame  of  hydrogen.  It  will  be  seen  that  drops  of  water 
are  condensed  upon  it. 

Hydrogen  combines  with  many  other  elements  besides 
oxygen,  and  forms  some  of  the  most  important  and  in- 
teresting compounds,  such  as  hydrochloric  acid,  HC1 ; 
sulphuretted  hydrogen  or  hydrogen  sulphide,  H2S ; 
ammonia,  NH3;  marsh  gas,  CH4 ;  and  all  the  acids.  On 
account  of  its  affinity  for  oxygen  it  is  used  very  exten- 
sively in  the  laboratory  for  the  purpose  of  extracting 
oxygen  from  compounds  containing  it.  Thus,  when 
hydrogen  is  passed  over  heated  copper  oxide,  CuO,  it 
combines-  with  the  oxygen  to  form  water,  and  the  coppel 


COMPARISON  OF  OXYGEN  AND  HYDROGEN.         47 

is  left  in  the  free  or  uncombined  state.     The  reaction  is 
represented  thus : 

CuO  +  2H=H30  +  Cu. 

A  similar  reaction  takes  place  when  hydrogen  is  passed 
over  highly  heated  oxide  of  iron,  FeaO3 : 

Fe2O3  +  6H  =  2Fe  +  3H3O. 

The  removal  of  oxygen  from  a  compound  is  called 
reduction.  Reduction  is  therefore  plainly  the  opposite 
of  oxidation.  Any  substance  which  has  the  power  to 
abstract  oxygen  is  spoken  of  as  a  reducing  agent,  just  as 
any  substance  which  has  the  power  to  add  oxygen  to 
a  substance,  or  to  decompose  it  by  the  action  of  oxygen, 
is  called  an  oxidizing  agent. 

A  number  of  metals  have  the  power  to  absorb  a 
large  quantity  of  hydrogen  when  they  are  heated  to 
red  heat  in  the  gas.  This  phenomenon  is  shown  most 
strikingly  by  palladium,  which  under  the  most  favor- 
able conditions  takes  up  something  more  than  935  times 
its  own  volume  of  hydrogen.  The  gas  is  given  up  at 
elevated  temperature  in  a  vacuum.  When  it  absorbs 
hydrogen,  palladium  undergoes  marked  changes  in 
properties.  Its  volume  is  increased,  and  its  magnetic 
and  electric  properties  are  also  changed.  It  was  sug- 
gested by  Graham,  who  first  observed  this  phenomenon, 
that  the  hydrogen  held  in  combination  by  the  palladium 
is  something  quite  different  from  ordinary  hydrogen, 
and  that  it  must  have  some  properties  like  those  of  the 
so-called  metals.  He  therefore  called  the  combined 
hydrogen  hydrogenium. 

Comparison  of  Oxygen  and  Hydrogen. — Hydrogen  and 
oxygen  are  different  kinds  of  matter,  just  as  heat  and 
•electricity  are  different  kinds  of  energy.  Heat  can  be 
•converted  into  electrical  energy,  and  electrical  energy 
into  heat,  but  one  element  cannot  by  any  means  known 
to  us  be  converted  into  another.  They  are  apparently 
entirely  independent  of  each  other.  The  question  will 
therefore  suggest  itself,  whether,  in  spite  of  their  ap- 


48  INORGANIC  CHEMISTRY. 

parent  independence,  there  is  not  some  relation  be- 
tween the  different  elements  which  reveals  itself  by 
similarity  in  properties  ?  It  will  be  found  that  the  ele- 
ments can  be  separated  into  groups  or  families  accord- 
ing to  their  properties.  There  are  some  elements,  for 
example,  which  in  their  chemical  conduct  resemble  oxy- 
gen markedly.  These  elements  constitute  the  oxygen 
family.  So  far  as  hydrogen  is  concerned,  however,  it 
stands  by  itself.  There  is  no  other  element  which  con- 
ducts itself  like  it.  If  we  compare  it  with  oxygen,  we 
find  very  few  facts  which  indicate  any  analogy  between 
the  two  elements.  In  their  physical  properties  they 
are,  to  be  sure,  similar.  Both  are  colorless,  inodorous, 
tasteless  gases.  On  the  other  hand,  oxygen  combines 
readily  with  a  large  number  of  substances  with  which 
hydrogen  does  not  combine.  Oxygen,  as  we  have  seen, 
combines  easily  with  carbon,  sulphur,  phosphorus,  and 
iron.  It  is  a  difficult  matter  to  get  any  of  these  elements 
to  combine  directly  with  hydrogen.  Further  than  this, 
substances  which  combine  readily  with  hydrogen  do  not 
combine  readily  with  oxygen.  The  two  elements  ex- 
hibit opposite  chemical  properties.  What  one  can  do 
the  other  cannot  do.  This  oppositeness  of  properties  is 
favorable  to  combination  ;  for  not  only  do  hydrogen  and 
oxygen  combine  with  great  ease  under  proper  condi- 
tions, but,  as  we  shall  see  later,  it  is  a  general  rule  that 
elements  of  like  properties  do  not  readily  combine  with 
one  another,  while  elements  of  unlike  properties  do 
readily  combine. 


CHAPTER  IV. 

STUDY  OF  THE  ACTION  OF  HYDROGEN  ON  OXYGEN. 

Burning  of  Hydrogen. — Attention  has  already  been 
called  to  the  fact  that,  when  hydrogen  burns,  water  is 
formed.  It  is  now  necessary  that  this  reaction  should 
be  studied  more  thoroughly  with  the  view  of  discover- 
ing, as  far  as  possible,  exactly  what  takes  place.  One 
of  the  first  questions  to  be  answered  is  what  relation 
exists  between  the  weights  of  the  hydrogen  burned,  the 
oxygen  used  up  in  the  burning  process,  and  the  water 
formed  ?  But  to  weigh  gases  accurately  and  to  collect 
small  quantities  of  water  and  weigh  it  are  by  no  means 
simple  operations,  and  a  great  deal  of  work  has  been 
done  upon  the  problem  under  consideration.  No  good 
method  has  been  devised  for  the  quantitative  study  of 
the  combination  of  hydrogen  and  oxygen  by  ordinary 
combustion.  On  the  other  hand,  very  accurate  experi- 
ments on  the  subject  have  been  made  in  three  other 
ways,  a  brief  account  of  which  will  now  be  given. 

Method  of  Dumas. — The  first  accurate  experiments 
on  the  combustion  of  hydrogen  are  those  of  Dumas. 
The  method  employed  by  this  chemist  was  as  follows : 
He  passed  carefully  purified  hydrogen  over  heated 
copper  oxide  and  collected  the  water  formed.  The 
reaction  involved  is  that  represented  by  the  equation 

CuO  +  2H  =  H2O  +  Cu. 

The  weight  of  the  oxygen  that  entered  into  combi- 
nation with  hydrogen  was  obtained  by  weighing  the 
vessel  containing  the  copper  oxide  before  and  after  the 
experiment.  The  loss  in  weight  represented  the  weight 
of  the  oxygen  which  had  been  abstracted.  The  water 
was  collected  by  passing  the  gases  formed  through  an 
empty  glass  vessel  in  which  most  of  the  water  was  con- 
densed, and  then  through  tubes  containing  molten  caus- 
tic potash  and  phosphorus  pentoxide,  substances  which 

(49) 


50  INORGANIC  CHEMISTRY. 

have  a  marked  power  to  absorb  water  and  hold  it  in 
combination.  In  nineteen  experiments  he  obtained  a 
total  amount  of  water  weighing  945.439  grams,  and  the 
total  amount  of  oxygen  used  in  the  formation  of  this 
amount  of  water  was  840.161  grams.  According  to  these 
results  the  ratio  between  hydrogen  and  oxygen  in  water 
expressed  in  percentages  is  : 

Oxygen 88.864 

Hydrogen 11.136 

100.000 

Morley  with  great  pains  proceeded  as  follows:  He 
weighed  comparatively  large  volumes  of  oxygen  in  the 
form  of  gas ;  he  then  weighed  hydrogen  after  it  had  been 
absorbed  by  palladium.  The  hydrogen  being  expelled 
from  the  palladium,  it  was  brought  together  with  the 
oxygen,  with  which  it  was  caused  to  combine  by  means 
of  electric  sparks.  The  water  thus  formed  was  carefully 
collected  and  weighed.  Twelve  experiments  gave  results 
that  agreed  closely  with  one  another.  The  mean  of  these 
shows  that  hydrogen  and  oxygen  combine  in  the  ratio  of 
1  part  of  hydrogen  to  15.8792  parts  of  oxygen.  Dumas' 
result  expressed  in  the  same  way  is  1 : 15.961. 

Eudiometric  Method. — When  hydrogen  and  oxygen  are 
mixed  at  ordinary  temperatures  no  chemical  change 
takes  place,  and  the  two  gases  may  be  left  in  contact 
with  each  other  indefinitely  without  chemical  action. 
If,  however,  a  spark  is  brought  in  contact  with  the 
mixture  violent  action  takes  place,  accompanied  by  a 
flame  and  explosion.  The  action  consists  in  the  sudden 
combination  of  the  two  gases  to  form  water.  It  may  be 
illustrated  in  a  number  of  ways  :  most  simply  by  filling 
soap-bubbles  with  the  mixture  of  the  two  gases  and  ap- 
plying a  flame  to  them.  The  explosion  which  ensues 
is  harmless.  Plainly,  to  study  the  combination  of 
hydrogen  and  oxygen  by  exploding  a  mixture  of  the 
gases  will  require  special  precautions.  It  can  be  carried 
out  by  the  aid  of  the  eudiometer  (from  evdia,  good  air, 
and  jteTpor,  a  measure,  an  instrument  for  determin- 


HYDROGEN  AND  OXYGEN— EUDIOMETRIC  METHOD.    51 

ing  the  purity  of  air).  The  eudiometer  is  simply  a 
tube  graduated  in  millimeters  and  having  two  small 
platinum  wires  passed  through  it  at  the  closed  end, 
nearly  meeting  inside  and  ending  in  loops  outside,  as 
shown  in  Fig.  1.  The  eudiometer  is  filled  with  mercury, 
inverted  in  a  mercury  trough,  and  held  in  an  upright 
position  by  means  of  proper  clamps.  For  the  purpose 
of  the  experiment  a  quantity  of  hydrogen  is  passed  up 
into  the  tube,  and  its  volume  accurately  measured. 
About  half  this  volume  of  oxygen  is  then  introduced 
and  the  volume  again  accurately  determined,  and 
after  the  mixture  has  been  allowed  to  stand  for  a 
few  minutes  a  spark  is  passed  between  the  wires  in 
the  eudiometer  by  connecting  the  loops  with  the  poles 


FIG.  1. 


of  a  small  induction  coil  or  with  a  Leyden  jar.  Under 
these  circumstances  the  explosion  takes  place  noise- 
lessly and  with  little  or  no  danger.  If  the  interior  of 
the  tube  was  dry  before  the  explosion,  it  will  be  seen  to 
be  moist  afterwards,  and  a  marked  decrease  in  the  vol- 
ume of  the  gases  is  also  observed.  That  water  is  the 
product  of  the  action  has  been  proved  beyond  any  possi- 
bility of  a  doubt,  over  and  over  again.  As  the  liquid 
water  which  is  formed  occupies  an  almost  inappreciable 
volume  as  compared  with  the  volume  of  the  gases  which 
combine,  the  decrease  in  volume  represents  the  total 
volume  of  hydrogen  and  oxygen  which  have  combined. 
Now,  if  the  experiment  is  performed  with  the  two  gases 
in  different  proportions,  it  will  be  found  that  only  when 
they  are  mixed  in  the  proportion  of  2-  volumes  of  hy- 
drogen to  1  volume  of  oxygen  do  they  completely  dis- 
appear in  the  explosion.  If  there  is  a  larger  proportion 
of  hydrogen  present,  the  excess  is  left  over;  and  the 
same  is  true  of  the  oxygen.  It  will  thus  be  seen  that 
when  hydrogen  and  oxygen  combine  to  form  water,  they 
do  so  in  the  proportion  of  2  volumes  of  hydrogen  to 
1  volume  of  oxygen  or  more  accurately  2.0008  to  1. 


52  INORGANIC  CHEMISTRY. 

Calculation  of  the  Results  Obtained  «in  Exploding  Mix- 
tures of  Hydrogen  and  Oxygen. — Having  determined  that 
whenever  hydrogen  and  oxygen  combine,  they  do  so  in 
the  proportion  1  volume  oxygen  to  2  volumes  hydrogen, 
and  that  when  they  combine  the  volume  of  liquid  water 
formed  measures  so  little  as  to  amount  to  nothing  in  the 
measurements,  we  know  that  whenever  a  mixture  of  hy- 
drogen and  oxygen  is  exploded,  no  matter  in  what  propor- 
tions they  may  be  present,  the  volume  of  gas  which  disap- 
pears as  such  consisted  of  2  volumes  of  hydrogen  and  1 
volume  of  oxygen,  or,  in  other  words,  one-third  of  the 
volume  which  disappears  was  oxygen  and  two-thirds 
hydrogen.  Take  this  example  :  A  quantity  of  hydrogen 
corresponding  to  60  cc.  under  standard  conditions  is  in- 
troduced into  a  eudiometer ;  40  cc.  of  oxygen  are  added* 
"What  contraction  will  there  be  on  exploding  the  mixture  ? 
Plainly  the  60  cc.  of  hydrogen  will  combine  with  30  cc. 
of  oxygen.  The  90  cc.  of  gas  will  disappear,  and  10  cc» 
of  oxygen  will  remain  uncombined.  From  a  total  vol- 
ume of  100  cc.,  therefore,  we  get  a  contraction  to  10  cc. 
One-third  of  the  contraction  represents  the  oxygen  and 
two-thirds  the  hydrogen. 

Determination  of  the  Volume  of  Water  Vapor  formed 
by  Union  of  Definite  Volumes  of  Hydrogen  and  Oxygen. 
— The  experiments  which  have  been  described  enable 
us  to  draw  the  conclusion  that  hydrogen  and  oxygen 
combine  in  certain  proportions  by  volume  and  by  weight, 
and  that  a  definite  weight  of  water  is  formed ;  further, 
that  the  volume  of  liquid  water  formed  when  the  two 
gases  combine  is  inappreciable  as  compared  with  that  of 
the  gases.  The  question  remains  to  be  answered,  what 
relation  exists  between  the  volumes  of  the  combining 
gases  and  that  of  the  water  in  the  form  of  vapor  ?  This 
can  be  determined  by  causing  the  gases  to  combine  in  a 
eudiometer  at  a  temperature  sufficiently  high  to  keep 
the  water  in  the  form  of  vapor.  The  simplest  arrange- 
ment for  accomplishing  this  is  that  shown  in  Fig.  2.  A 
long  eudiometer,  the  upper  half  of  which  is  divided  into 
three  equal  divisions  marked  on  the  outside,  is  filled 
with  mercury  and  inverted  in  a  bath  of  mercury.  It  is 


HYDROGEN  AND  OXYGEN— WATER  VAPOR. 


53 


then  surrounded  by  a  large  tube  or  jacket  arranged  as 
shown  in  Fig.  2.  This  is  connected  through  the  cork 
at  the  lower  end  with  a  vessel  from  which  a  current  of 
steam  can  be  obtained.  The  steam  is  passed  through 
the  jacket  until  the  temperature  of  the  mercury  has 
reached  that  of  the  steam.  A  mixture  of  hydrogen  and 
oxygen  in  the  proportions  in  which  they  combine,  viz., 
2  of  hydrogen  to  1  of  oxygen,  is  then  introduced  into 
the  eudiometer  so  that  it  is  filled  to  the  third  mark. 
This  must  be  somewhat  above  the  level  of  the  mercury 
in  the  bath,  so  that  the  gases  in  the  eudiometer  shall  be 
under  diminished  pressure.  On  now  passing  the  spark 
the  gases  unite  and  the  water  which 
is  formed  remains  in  the  form  of 
vapor,  as  the  temperature  inside 
the  eudiometer  is  nearly  that  of 
boiling  water,  and  the  vapor  is  un- 
der diminished  pressure.  It  is 
found  that  the  volume  of  the  water 
vapor  is  less  than  that  of  the  gases 
introduced  into  the  eudiometer. 
In  order  to  secure  the  same  press- 
ure as  that  under  which  the  gases 
were  measured,  the  eudiometer 
must  be  lowered  until  the  height 
of  the  mercury  column  in  it  is  the 
same  as  it  was  before  the  explosion. 
On  now  measuring  the  volume  of 
water  vapor,  it  will  be  found  to  be 
two-thirds  that  occupied  by  the 
uncombined  gases.  Therefore,  2 
volumes  of  hydrogen  combine  with 
1  volume  of  oxygen  and  form  2 
volumes  of  water  vapor.  It  is  an 

interesting  fact  that  these  simple  relations  exist  between 
the  volumes  of  the  combining  gases  and  the  volume  of 
the  product.  We  shall  see  that  similar  relations  hold 
good  in  the  case  of  other  gases  ;  and  the  following  gen- 
eral statement  is  based  upon  a  great  deal  of  careful 
study  : 


FIG.  2. 


54  INORGANIC  CHEMISTRY. 

When  two  or  more  gaseous  substances  combine  to  form  a 
gaseous  compound,  the  volumes  of  the  individual  constituents 
as  well  as  their  sum  bear  a  simple  relation  to  the  volume  of 
the  compound. 

This  is  known  as  the  law  of  combination  by  volume. 
As  will  be  seen  farther  on,  it  has  a  most  important  bear- 
ing upon  some  of  the  fundamental  ideas  held  in  regard 
to  the  constitution  of  matter. 

Heat  Evolved  in  the  Union  of  Hydrogen  and  Oxygen. — To 
get  as  complete  a  knowledge  as  possible  of  the  reaction 
which  takes  place  between  hydrogen  and  oxygen  we  have 
still  to  determine  the  amount  of  heat  evolved.  The  heat 
evolved  in  burning  a  gram  of  hydrogen  can  be  deter- 
mined, and  from  this  we  can  calculate  the  heat  of  forma- 
tion of  water  which,  according  to  what  was  said  on  page 
37,  is  the  amount  of  heat  evolved  by  the  combination 
of  2  grams  of  hydrogen  with  15.88  grams  of  oxygen.  We 
are  to  determine  the  value  of  x  in  the  equation 

[H2,  O]  =  x  cal, 

which  expresses  in  thermochemical  language  the  fact  that 
when  2  grams  of  hydrogen  combine  with  15.88  grams  of 
oxygen  x  calories  of  heat  are  evolved. 

The  determination  is  made  by  burning  a  known  weight 
of  hydrogen  in  a  vessel  surrounded  by  water  and  arranged 
in  such  a  way  that  all  the  heat  is  absorbed  by  the  water. 
Experiment  has  shown  that  when  1  gram  of  hydrogen  is 
burned  34,180  calories  are  evolved.  Or, 

[H2,  O]  =  68,360  cal. 

No  other  substance  gives  as  much  heat  as  this  in  propor- 
tion to  the  weight  used. 

Applications  of  the  Heat  formed  by  the  Combination  of 
Hydrogen  and  Oxygen. — To  burn  hydrogen  in  the  air  is, 
as  we  have  seen,  a  simple  matter,  but  to  burn  it  in  oxy- 
gen requires  a  special  apparatus  to  prevent  the  mixing 
of  the  gases  before  they  reach  the  end  of  the  tube  where 
the  combustion  takes  place.  The  oxyhydrogen  bloiv-pipe 
answers  this  purpose.  This  may  be  constructed  in  sev- 


OXYHYDROGEN  BLO  W-PIPE— OXYHYDROGEN  LIGHT.     55 

eral  ways,  but  the  simplest  is  that  represented  in  Fig.  3. 
It  consists  of  a  tube  through  which  a  smaller  tube  passes. 
The  hydrogen  is  admitted  through  a  and  the  oxygen 
through  b.  It  will  be  seen  that  they  come  together  only 
at  the  end  of  the  tube.  The  hydrogen  is  first  passed 
through  and  lighted ;  then  the  oxygen  is  passed  through 
slowly,  the  pressure  being  increased  until  the  flame  ap- 
pears thin  and  straight.  It  gives  very  little  light  but  is 
intensely  hot.  Iron  wire,  steel,  copper,  zinc,  and  other 


a  FIG.  3. 

metals  burn  in  the  flame  with  ease.  Platinum  vessels 
are  made  by  melting  the  platinum  by  means  of  the  oxy- 
hydrogen  flame. 

Oxyhydrogen  Light. — When  the  oxyhydrogen  flame  is 
allowed  to  play  upon  some  substance  which  it  cannot 
melt  or  burn,  the  substance  becomes  heated  so  high  that 
it  gives  off  an  intense  light.  The  substance  commonly 
used  for  this  purpose  is  quicklime.  Hence  the  light  is 
often  called  the  lime-light.  It  is  also  known  as  the 
Drummond  light. 

The  hydrogen  is  first  allowed  to  pass  through  the  stop- 
cock, and  lighted,  when  the  oxygen  is  admitted.  The 
flame  plays  against  the  piece  of  lime,  and  from  this  the 
light  is  given  off  when  it  has  acquired  a  high  tempera- 
ture. Coal-gas  may  be  used  instead  of  hydrogen,  and  it 
is  generally  used.  Cylinders  of  compressed  coal-gas  and 
of  oxygen  can  be  bought,  and  the  gases  so  prepared  are 
used  for  the  purpose  of  projecting  pictures  upon  screens 
in  illustrating  lectures  and  for  other  similar  purposes. 

Velocity  of  Combination  of  a  Mixture  of  Hydrogen  and 
Oxygen. — When  a  mixture  of  hydrogen  and  oxygen  ex- 
plodes, the  action  appears  to  take  place  instantaneously 
throughout  the  mass.  Whether  this  is  really  so  or  not 
can  be  determined  only  by  experiment.  The  action  cer- 


56  INORGANIC  CHEMISTRY. 

tainly  takes  place  with  great  rapidity,  and  a  special 
apparatus  is  necessary  in  order  to  study  the  rate  of 
transmission.  This  subject  has  been  studied  by  ex- 
ploding the  mixture  in  long  tubes  arranged  with  little 
movable  pistons  at  various  distances.  When  the  explo- 
sion reached  these  points  the  fact  was  indicated  by 
motion  of  the  pistons.  The  result  showed  that  the 
action  is  not  instantaneous  though  extremely  rapid.  The 
rate  of  transmission  is  about  2500  metres  per  second. 

Summary. — In  our  study  of  the  action  of  hydrogen  on 
oxygen,  we  have  learned :  (1)  the  relations  between  the 
weights  of  the  two  gases  which  act  upon  each  other  ;  (2) 
the  relations  between  the  volumes  of  the  combining 
gases  ;  (3)  the  relations  between  the  volumes  of  the  com- 
bining gases  and  that  of  the  water  vapor  formed  ;  (4)  the 
amount  of  heat  evolved  when  a  given  weight  of  hydrogen 
combines  with  oxygen ;  and  (5)  that  the  act  of  combina- 
tion of  the  two  gases  does  not  take  place  instantaneously 
though  with  great  rapidity.  It  remains  for  us  to  study 
more  carefully  the  product  formed.  This  is  water. 


CHAPTER  V. 

WATER. 

Historical. — Water  was  long  considered  an  elementary 
substance  until,  towards  the  end  of  the  last  century,  the 
discovery  of  hydrogen  and  oxygen,  and  of  the  nature  of 
combustion,  led  to  the  discovery  of  its  composition. 

Occurrence. — Besides  the  form  in  which  water  occurs 
in  such  enormous  quantities  in  the  earth,  it  also  occurs 
in  forms  and  conditions  which  prevent  its  immediate 
recognition.  Thus  all  living  things  contain  a  large  pro- 
portion of  water,  which  can  be  driven  off  by  heat.  If  a 
piece  of  wood  or  of  meat  is  heated,  liquids  pass  off,  and 
by  purification  these  can  be  shown  to  consist  mainly  of 
water.  The  proportion  of  water  in  animal  and  vegetable 
substances  is  very  great.  If  the  body  of  a  man  weighing 
150  pounds  were  to  be  put  in  an  oven  and  thoroughly 
dried,  there  would  be  only  about  40  pounds  of  solid  mat- 
ter left,  most  of  the  rest  being  water. 

Water  also  occurs  in  another  form  in  which  it  does  not 
easily  reveal  its  presence.  This  is  as  water  of  crystalliza- 
tion. Many  chemical  compounds  found  in  nature  and 
manufactured  are  found  to  give  off  water  when  heated. 
If,  for  example,  the  zinc  sulphate  formed  in  the  prep- 
aration of  hydrogen  from  zinc  and  sulphuric  acid  is 
dried  by  exposure  to  the  air  or  by  pressing  between  lay- 
ers of  filter-paper,  it  will  be  found  that,  when  heated  in 
a  dry  tube,  it  gives  off  water,  and  at  the  same  time 
changes  its  appearance.  The  same  is  true  of  gypsum 
which  is  found  in  nature,  and  of  copper  sulphate  or  blue 
vitriol.  In  this  last  case  the  loss  of  water  is  accompanied 
by  a  loss  of  color.  After  all  the  water  is  driven  off,  the 
powder  left  behind  is  white. 

Many  compounds  when  deposited  from  solutions  in 
water  in  the  form  of  crystals  combine  with  definite  quan- 

(57) 


58  INORGANIC  CHEMISTRY. 

titles  of  water.  This  water  is  not  present  as  such,  but  is 
held  in  chemical  combination.  Hence  the  substance  does 
not  appear  moist,  though  it  may  contain  more  than  half 
its  weight  of  water.  This  water  of  crystallization  is,  in 
some  way  which  is  not  understood,  essential  to  the  form 
of  the  crystals.  If  it  is  driven  off,  the  crystals  generally 
crumble  to  pieces.  Some  compounds  combine  with  dif- 
ferent quantities  of  water  under  different  circumstances, 
the  form  of  the  crystals  varying  with  the  quantity  of 
water  held  in  combination. 

Compounds  differ  greatly  as  regards  the  ease  with 
which  they  give  up  water  of  crystallization.  In  general, 
it  is  given  off  when  the  compound  containing  it  is  heated 
up  to  the  temperature  of  boiling  water.  But  some  com- 
pounds give  it  up  by  simple  contact  with  the  air.  This 
is  true  of  sodium  sulphate  or  Glauber's  salt,  which  con- 
tains a  quantity  of  water  of  crystallization  represented  by 
the  formula  Na2SO4.  10H2O.  If  some  of  the  crystals  are 
allowed  to  lie  exposed  to  the  air  they  undergo  a  marked 
change  in  the  course  of  an  hour  or  two.  They  lose 
their  lustre  and  gradually  crumble  to  pieces.  Sub- 
stances which  lose  their  water  of  crystallization  by  sim- 
ple contact  with  the  air  are  said  to  be  efflorescent. 

Some  compounds  if  deprived  of  their  water  of  crystal- 
lization will  take  it  up  again  when  allowed  to  lie  in  an 
atmosphere  containing  moisture.  As  the  air  always  con- 
tains moisture,  it  is  only  necessary  to  expose  such  com- 
pounds to  the  air  in  order  to  notice  the  phenomenon. 
It  is  well  shown  by  the  compound  calcium  chloride, 
CaCl2.  This  substance  has  a  remarkable  power  of  at- 
tracting water  and  holding  it  in  combination.  If  a  few 
pieces  are  exposed  to  the  air  it  will  be  noticed  that  they 
soon  have  a  moist  appearance,  and  if  they  are  allowed  to 
lie  long  enough  they  will  dissolve  in  the  water  which  is 
absorbed  from  the  air.  Substances  which  absorb  water 
from  the  air  are  said  to  be  deliquescent. 

Formation  of  Water  and  Proofs  of  its  Composition. — If 
we  had  not  already  learned,  in  studying  the  action  of  hy- 
drogen upon  oxygen,  that  water  is  composed  of  these 
two  gases,  we  should  first  subject  it  to  analysis.  For 


FORMATION  OF  WATER.  59 

this  purpose  we  should  have  to  bring  it  under  the  influ- 
ence of  a  number  of  reagents  and  study  its  conduct.  If 
we  should  pass  an  electric  current  through  it  in  the  proper 
way,  we  should  observe  that  a  gas  rises  from  each  pole. 
By  placing  each  pole  under  the  mouth  of  an  inverted 
tube  filled  with  water  the  gases  are  easily  collected. 
When  one  of  the  tubes  has  become  full  of  gas  the  other 
one  will  be  only  half  full.  An  examination  will  show 
that  the  gas  in  the  tube  which  is  filled  is  hydrogen, 
whereas  that  in  the  one  which  is  half  filled  is  oxygen. 
By  decomposition  of  water  by  means  of  the  electric  cur- 
rent, then,  there  are  obtained  two  volumes  of  hydrogen 
for  each  volume  of  oxygen.  We  already  know  the  rela- 
tive weights  of  equal  volumes  of  the  two  gases,  so  that 
we  can  easily  calculate  the  relative  weights  of  the  gases 
obtained  in  the  experiment.  The  ratio  of  the  weights 
of  equal  volumes  of  hydrogen  and  oxygen  is  1 : 15.88. 
Therefore,  if  we  have  2  volumes  of  hydrogen  combined 
with  1  volume  of  oxygen,  the  ratio  between  the  weights 
is  2  :  15.88  or  1  :  7.94  Although  we  know  from  the  ex- 
periment referred  to  that  hydrogen  and  oxygen  are  ob- 
tained from  water  in  certain  proportions,  it  does  not  fol- 
low that  this  is  the  composition  of  water.  For  it  may  be 
that  other  elements  besides  hydrogen  and  oxygen  are 
contained  in  it,  and  it  may  be  also  that  all  the  hydrogen 
and  oxygen  are  not  set  free  by  the  action  of  the  electric 
current.  We  might  determine  whether  either  of  these 
possibilities  is  true  or  not  by  decomposing  a  weighed 
quantity  of  water,  and  weighing  the  hydrogen  and  oxygen 
obtained  from  it.  If  we  should  find  that  the  sum  of  the 
weights  of  hydrogen  and  oxygen  is  equal  to  the  weight 
of  the  water  decomposed,  this  fact  would  be  evidence 
that  only  hydrogen  and  oxygen  are  contained  in  water, 
and  that  they  are  present  in  the  proportions  stated. 
The  same  thing  can  be  satisfactorily  proved  by  causing 
hydrogen  and  oxygen  to  combine,  or  by  effecting  the  syn- 
thesis of  water.  How  this  may  be  done  has  already  been 
pointed  out.  It  was  shown,  in  the  first  place,  that  by 
burning  hydrogen  in  oxygen  water  is  formed.  This 
proves  that  water  consists  of  hydrogen  and  oxygen,  but 


60  INORGANIC  CHEMISTRY. 

it  does  not  furnish  any  proof  as  to  the  relation  between 
the  quantities  of  the  gases  which  combine.  It  is  a  quali- 
tative synthesis.  Other  methods  were  described,  the  ob- 
ject of  which  was  to  show  in  what  proportion  by  weight 
and  by  volume  hydrogen  and  oxygen  combine  to  form 
water.  These  methods  are  examples  of  quantitative  syn- 
theses. The  results  proved  that  to  form  water  hydrogen 
and  oxygen  combine  in  the  proportion  of  1  volume  of 
oxygen  to  2  of  hydrogen ;  and  it  therefore  follows  that 
the  decomposition  of  water  which  is  effected  by  the  elec- 
tric current  is  complete. 

Properties  of  Water. — Pure  water  is  tasteless  and  in- 
odorous, and  in  small  quantities  colorless.  Thick 
layers  are,  however,  blue.  This  is  seen  by  filling  a 
long  tube  with  carefully  purified  water,  and  examining  it 
by  transmitted  light,  when  it  appears  blue.  Some 
mountain  lakes  also  have  a  marked  blue  color.  When 
cooled,  water  contracts  until  it  reaches  the  temperature 
of  4°.  At  this  point  it  has  its  maximum  density.  If 
cooled  below  this  it  expands,  and  the  specific  gravity  of 
ice  is  somewhat  less  than  that  of  water.  Hence  ice 
floats  on  water.  If  this  were  not  so  there  would  be 
great  danger  in  cold  climates  that  the  water  in  the 
streams  would  freeze  solid.  As  it  is,  the  lower  layers 
of  water  are  protected  by  the  ice  and  the  cold  water 
just  below  it,  which  are  poor  conductors  of  heat.  Water 
can  be  cooled  down  below  its  freezing  temperature,  or 
0°,  if  it  is  kept  perfectly  quiet,  protected  from  the  air,  or 
cooled  in  capillary  tubes.  Water  thus  cooled  down  will 
suddenly  solidify  when  disturbed,  and  then  its  tempera- 
ture rises  to  0°.  Water  boils  at  100°  under  760  mm.  pres- 
sure. Increased  pressure  raises  the  boiling  point,  and 
decreased  pressure  lowers  it. 

Chemical  Properties  of  Water. — Water  is  a  very  stable 
chemical  compound.  An  indication  of  the  attraction  ex- 
erted by  the  hydrogen  for  the  oxygen  is  given  in  the 
great  evolution  of  heat  when  the  two  combine.  In  order 
to  decompose  it  by  heat,  as  much  heat  must  be  added  as 
is  evolved  when  it  is  formed.  At  high  temperatures  it 
is  decomposed  into  its  elements.  The  decomposition 


CHEMICAL  PROPERTIES  OF  WATER.  61 

begins  at  1000°  and  is  half  complete  at  2500°.  This 
kind  of  gradual  decomposition  of  a  compound  by  heat 
is  called  dissociation.  It  is  a  common  phenomenon  in 
chemistry,  and  farther  on  we  shall  have  occasion  to  study 
it  more  fully.  As  has  been  seen,  water  is  decomposed 
completely  by  an  electric  current,  and  partly  by  contact 
with  sodium  and  potassium  at  ordinary  temperatures, 
and  by  iron  and  carbon  at  higher  temperatures.  It 
combines  directly  with  a  large  number  of  substances  in 
the  form  of  water  of  crystallization,  and  with  others  to 
form  definite  chemical  compounds  called  hydrates  or 
hydroxides.  Thus  the  oxides  of  potassium,  K2O,  and  of 
sodium,  Na2O,  combine  with  water  with  evolution  of 
much  heat  to  form  the  compounds  potassium  hydroxide, 
KOH,  and  sodium  hydroxide,  NaOH,  which,  it  will  be  re- 
membered, are  also  formed  by  the  action  of  the  elements 
potassium  and  sodium  on  water  with  liberation  of  hydro- 
gen. The  reactions  between  the  oxides  and  water  are 
represented  by  the  equations 

K2O   +  H20=2KOH; 
Na2O  +  H20  =  2NaOH. 

Similarly,  lime  or  calcium  oxide,  CaO,  acts  upon  water 
with  evolution  of  heat,  as  is  observed  in  the  process  of 
slaking.  The  change  is  like  that  which  takes  place  with 
potassium  and  sodium  oxides  ;  and  is  represented  thus  : 

CaO  +  H20  =  CaO2H2  [or  Ca(OH)J. 

The  product  represented  by  the  symbol  Ca(OH)2  is 
known  as  calcium  hydroxide  or  slaked  lime.  In  the 
same  way  barium  oxide,  BaO,  forms  barium  hydroxide  : 

BaO  +  H2O  =  BaOJI2  [or  Ba(OH)J. 

We  shall  meet  with  many  examples  of  this  kind  of 
action  in  our  study  of  chemical  reactions,  and  we  shall 
see  that  the  hydroxides  form  two  of  the  most  important 
classes  of  compounds,  known  as  acids  and  bases.  The 
hydroxides  of  potassium,  sodium,  calcium,  and  barium 
are,  for  example,  bases ;  while  certain  hydroxides  con- 


62  INORGANIC  CHEMISTRY. 

taining  sulphur,  nitrogen,  and  carbon  are  acids,  —  such  as 
sulphuric  acid  SO2(OH)2,  nitric  acid  NO2(OH),  and  car- 
bonic acid,  CO(OH)2.  It  is  not  believed  that  water  as 
such  is  contained  in  these  hydroxides.  Nevertheless, 
when  heated  many  of  them  give  off  water.  Thus,  when 
heated  to  a  red  heat,  calcium  hydroxide  is  decomposed 
into  the  oxide  and  water  according  to  the  equation 


Many  substances  which  contain  hydrogen  and  oxygen 
act  in  the  same  way.  This  is  due  to  the  great  sta- 
bility of  water  even  at  elevated  temperatures.  As  the 
temperature  becomes  higher  and  higher  the  attraction 
between  the  constituents  of  the  compound  becomes 
weaker  and  weaker.  When  a  point  is  reached  at  which 
the  attraction  of  the  hydrogen  for  the  oxygen  is  greater 
than  that  required  to  hold  the  constituents  together,  a 
rearrangement  takes  place,  and  compounds  which  are 
stable  at  the  higher  temperature  are  formed. 

Water  as  a  Solvent.  —  With  a  great  many  substances 
water  forms  unstable  compounds,  the  nature  of  which 
cannot  at  present  be  explained.  These  unstable  com- 
pounds are  called  solutions.  It  is  known  that  many  solids, 
liquids,  and  gases  when  brought  into  water  disappear 
and  form  colorless  liquids  which  look  like  water.  Some 
give  colored  liquids  of  the  same  color  as  the  substance 
dissolved,  and  others  give  liquids  which  have  a  color 
quite  different  from  the  substance  dissolved.  On  the 
other  hand,  there  are  many  substances  which  do  not 
form  such  compounds  with  water  or  which,  as  we  say, 
are  insoluble  in  water.  In  a  solution  the  particles  of 
the  substance  dissolved  are  in  some  way  attracted  and 
held  in  combination  by  the  particles  of  the  liquid.  If  a 
very  small  quantity  of  substance  be  dissolved  in  a  large 
quantity  of  water  and  the  solution  thoroughly  stirred, 
the  dissolved  substance  is  uniformly  distributed  through- 
out the  liquid,  as  can  be  shown  by  refined  chemical  meth- 
ods. That  the  dissolved  substance  is  everywhere  present 
in  the  solution  can  be  shown  further  by  the  aid  of  certain 
dye-stuffs,  as,  for  example,  fuchsine.  A  drop  of  a  concen- 


WATER  AS  A  SOLVENT.  63 

trated  solution  of  the  substance  brought  into  many  gallons 
of  water  imparts  a  distinct  color  to  all  parts  of  the  liquid. 
An  experiment  of  this  kind  gives  some  idea  of  the 
extent  to  which  the  division  of  matter  can  be  carried. 
For  it  is  evident  that  in  each  drop  of  the  dilute  solution 
there  must  be  contained  some  of  the  coloring  matter, 
though  the  quantity  must  be  what  we  should  ordinarily 
speak  of  as  infinitesimal.  While  there  seems  to  be  no 
limit  to  the  extent  to  which  a  solution  can  be  diluted  and 
still  retain  the  dissolved  substance  uniformly  distributed 
through  its  mass,  there  is  a  limit  to  the  amount  of  every 
substance  that  can  be  brought  into  solution,  and  this 
varies  with  the  temperature,  and,  in  the  case  of  gases, 
with  the  pressure.  Some  substances  are  easily  soluble, 
others  are  difficultly  soluble.  When  the  solutions  are 
boiled  the  water  simply  passes  off  and  leaves  the  dis- 
solved substance  behind,  if  it  is  a  non-volatile  solid. 
If,  however,  the  substance  in  solution  is  a  liquid,  a  par- 
tial separation  will  take  place,  the  extent  of  the  separa- 
tion depending  largely  upon  the  difference  between  the 
boiling-points  of  the  water  and  the  other  liquid.  A  com- 
plete separation  of  two  liquids  by  boiling  is  difficult  and 
in  most  cases  impossible.  If,  finally,  the  substance  in 
solution  is  a  gas,  it  generally  passes  off  when  the  solution 
is  heated,  though  in  some  cases  water  is  given  off  leaving 
the  gas  in  solution,  which  of  course  then  becomes  more 
concentrated.  When  a  certain  concentration  is  reached  a 
solution  of  the  gas  passes  over.  It  is  probable  that  in  these 
cases  the  gas  is  in  a  condition  of  chemical  combination 
with  the  water.  Solutions,  in  general,  seem  to  differ  from 
true  chemical  compounds  in  some  important  particulars, 
and  also  from  mere  mechanical  mixtures.  Definiteness 
of  composition  appears  to  be  characteristic  of  chemical 
compounds,  or,  at  least,  it  is  characteristic  of  a  large  num- 
ber of  compounds  which  we  call  chemical  compounds. 
But  solutions  have  no  definite  composition.  We  can  dis- 
solve any  quantity  of  a  substance  from  the  minutest  par- 
ticle to  a  certain  fixed  quantity,  and  the  solutions  formed 
are  uniform  and  appear  to  be  just  as  truly  solutions  as 
that  which  contains  the  largest  quantity  which  can  be 


64  INORGANIC  CHEMISTRY. 

held  in  combination.  On  the  other  hand,  in  a  mere  me- 
chanical mixture  the  constituents  may  be  present  in  all 
proportions,  while  this  is  not  true  of  solutions.  The  sub- 
ject of  solution  is  under  investigation,  and  it  has  been 
made  highly  probable  that  some  substances,  when  dis- 
solved, are  partly  or  wholly  broken  down  into  smaller 
parts  charged  with  positive  and  negative  electricity  re- 
spectively. These  smaller  parts  are  called  ions. 

Solution  as  an  Aid  to  Chemical  Action. — When  it  is 
desired  to  secure  the  chemical  action  of  one  solid  sub- 
stance upon  another,  it  is  generally  necessary  to  bring 
them  together  in  solution.  One  reason  why  they  do 
not  act  readily  when  mixed  in  the  solid  condition  is  to 
be  found  in  the  fact  that,  under  these  circumstances, 
their  particles  remain  separated  by  sensible  distances, 
no  matter  how  finely  the  mixture  may  be  powdered. 
If,  however,  the  substances  are  dissolved,  and  the 
solutions  poured  together,  the  particles  of  the  liquid 
move  so  freely  among  one  another  that  they  come 
in  intimate  contact,  thus  facilitating  chemical  action. 
Many  substances  which  do  not  act  upon  one  another  at 
all  when  brought  together  in  dry  condition  act  readily 
when  brought  together  in  solution.  It  is  believed  that 
this  is  due  principally  to  the  splitting  of  the  compounds 
into  ions  by  the  action  of  the  water.  These  ions  being 
free  are  capable  of  acting  upon  other  ions  which  may 
be  brought  into  the  same  solution.  This  idea  will  be 
developed  farther  on  in  other  connections.  Although 
it  is  highly  probable,  £hen,  that  when  a  reaction  takes 
place  in  a  water  solution  the  water  itself  plays  a  very 
important  part,  the  reaction  is  generally  represented 
by  an  equation  in  which  the  water  does  not  appear.  Of 
course,  such  an  equation  is  imperfect,  but  it  answers  cer- 
tain purposes  quite  satisfactorily,  and  may  be  used  with- 
out danger  of  confusion.  Thus,  when  hydrochloric  acid 
acts  upon  zinc,  hydrogen  is  liberated  and  zinc  chloride 
is  formed.  What  we  call  hydrochloric  acid  in  the  labo- 
ratory is  the  liquid  which  is  formed  by  the  absorption  of 
hydrochloric  acid  gas,  HC1,  by  water.  When  this  liquid 
is  used,  however,  the  chemical  act  which  makes  itself 


NATURAL   WATEES.  65 

known  to  us  is  that  which  gives  hydrogen  gas  and  zinc 
chloride  in  solution.  Apparently  this  reaction  is  inde- 
pendent of  the  water,  and  it  may  be  represented  thus  : 

Zn  +  2HC1  =  ZnCla  +  2H. 

Sometimes  such  reactions  are  written  as  follows  in  order 
to  express  the  fact  that  water  is  present,  though,  as  will 
be  observed,  no  attempt  is  made  to  tell  what  part  the 
water  plays : 

Zn  +  2HC1  +  Aq    =  ZnCl2  +  2H  +  Aq  ;  or 
Zn  +  2HC1  +  H2O  =  ZnCl3  +  2H  +  H2O. 

There  is  no  objection  to  this,  certainly,  but  it  is  ques- 
tionable whether,  considering  the  purposes  for  which 
chemical  equations  are  used,  this  increases  their  value. 

Natural  Waters. — All  water  found  in  nature  is  more  or 
less  impure  or,  in  other  words,  contains  something  in  so- 
lution. In  the  first  place,  waters  which  are  exposed  to 
the  air  dissolve  some  of  the  gases  of  which  the  air  is 
composed,  as  oxygen,  nitrogen,  and  carbon  dioxide. 
Again,  natural  waters  necessarily  are  in  contact  with 
the  earth,  they  always  dissolve  some  of  the  earthy 
substances ;  and,  finally,  many  waters  come  in  contact 
with  animal  and  vegetable  substances  and  dissolve  some- 
thing from  these.  The  water  which  is  carried  up  as 
vapor  from  the  surfaces  of  natural  bodies  of  water  is 
approximately  pure.  When  this  is  precipitated  as  rain  it 
dissolves  certain  substances  from  the  air,  and  the  first  rain 
that  falls  during  a  storm  is  always  more  or  less  contami- 
nated. In  a  short  time,  however,  the  air  becomes  washed 
and  the  rain  which  falls  thereafter  is  approximately  pure 
water.  If  it  remains  in  contact  with  insoluble  rocks,  as, 
for  example,  quartzite  or  sandstone,  it  remains  pure,  and 
mountain-streams  which  flow  over  sandstone  beds  are,  in 
general,  the  purest.  Water  which  flows  over  limestone 
dissolves  some  of  this  and  becomes  "  hard."  A  similar 
change  is  brought  about  in  water  by  contact  with  gyp- 
sum and  magnesium  sulphate.  The  condition  of  hard- 
ness will  be  taken  up  more  fully  under  calcium  and 
magnesium  compounds.  The  many  varieties  of  mineral 


66  INORGANIC  CHEMISTRY. 

springs  have  their  origin  in  the  presence  in  the  earth  of 
certain  substances  which  are  soluble  in  water.  Among 
those  most  frequently  met  with  in  solution  in  natural 
waters  are  carbonic  acid,  sodium  carbonate,  sodium  sul- 
phate or  Glauber's  salt,  sodium  chloride  or  common  salt, 
magnesium  sulphate,  iron  carbonate,  and  sulphuretted 
hydrogen.  Effervescent  waters  are  those  which  contain 
a  large  quantity  of  carbonic  acid  in  solution  and  give 
off  carbon  dioxide  gas  when  exposed  to  the  air.  Chalyb- 
eate waters  are  those  which  contain  some  compound  of 
iron  in  solution  ;  sulphur  waters  contain  the  gas,  sulphu- 
retted hydrogen.  Common  salt  occurs  in  large  quantities 
in  different  parts  of  the  earth.  As  it  is  easily  soluble  in 
water,  many  streams  contain  it ;  and  as  most  streams  find 
their  way  to  the  ocean,  we  see  one  reason  why  the  water 
of  the  ocean  should  be  salt. 

As  streams  approach  the  habitation  of  man  they  are 
subjected  to  a  serious  cause  of  contamination.  The 
drainage  from  the  neighborhood  of  human  dwellings  is 
very  apt  to  find  its  way  into  a  near  stream.  The  sub- 
stances thus  carried  into  the  stream  undergo  decompo- 
sition and  give  rise  to  the  formation  of  larger  or  smaller 
quantities  of  new  products  some  of  which  have  the  power 
to  produce  disturbances  when  taken  into  the  system, 
and  others  to  produce  disease.  This  condition  of  things 
is  most  strikingly  illustrated  by  the  case  of  a  large  town 
situated  on  the  banks  of  a  river.  It  frequently  happens 
that  the  water  of  the  river  is  used  for  drinking  purposes, 
and  it  also  frequently  happens  that  the  water  is  contami- 
nated by  drainage.  Biver  water  when  once  contaminated 
by  drainage  tends  to  become  pure  again  by  contact  with 
the  air,  the  change  consisting  largely  in  the  slow  oxida- 
tion of  the  substances  which  are  of  animal  or  vegetable 
origin,  and  their  conversion  into  harmless  products.  If 
water  is  to  be  used  for  drinking  purposes,  however,  it  is 
not  well  to  rely  too  much  upon  this  process  of  purifica- 
tion. So  much  has  of  late  years  been  said  about  drinking- 
water  that  excessive  alarm  has  been  created,  and  water  is 
no  doubt  frequently  held  responsible  for  sickness  with 
which  it  has  nothing  to  do.  In  some  places  the  war  against 
the  water  supply  has  been  carried  so  far  that  those  who 


PURIFICATION  OF  WATER.  67 

can  afford  it  drink  only  artificially  purified  and  distilled 
water.  It  is  undoubtedly  well  to  be  cautious,  but  it 
is  possible  to  be  too  cautious. 

What  Constitutes  a  Bad  Drinking  Water. — A  good  drink- 
ing water  should  be  free  from  odor  and  taste  and  should 
not  contain  anything  which  can  act  injuriously  upon  the 
system.  It  is,  however,  difficult  to  decide  by  chemical 
means  whether  the  water  contains  anything  injurious  or 
not,  as  there  may  be  a  very  minute  quantity  of  an  ex- 
tremely injurious  substance,  for  example  a  disease  germ, 
present,  and  chemical  analysis  would  be  powerless  to  de- 
tect it.  On  the  other  hand,  water  which  is  very  consid- 
erably contaminated  by  sewage  may  be  harmless,  and 
yet  the  latter  might  be  pronounced  "  bad "  and  the 
former  "  good."  The  rule  generally  adopted  by  chemists 
in  dealing  with  water  is  to  pronounce  any  water  danger- 
ous which  is  contaminated  by  sewage.  Such  contamina- 
tion can  generally  be  detected  by  analysis  or  by  analysis 
and  inspection  of  the  sources. 

Purification  of  Water.— Impure  water  may  become  purer 
by  natural  methods  as  has  been  stated,  and  it  may  be 
rendered  fit  for  drinking  purposes  by  filtering  through 
such  substances  as  charcoal,  sand,  spongy  iron,  etc.  A 
filter,  no  matter  of  what  it  may  be  made,  will  not,  how- 
ever, remain  efficient  for  any  length  of  time,  as  the  sub- 
stances contained  in  the  impure  water  are  retained  by  it 
and,  after  a  time,  it  becomes  a  source  of  pollution  in- 
stead of  a  purifier.  For  refined  work  in  chemistry  pure 
water  is  prepared  by  distilling  natural  waters.  The 
process  of  distillation  consists  in  boiling  the  water  and 
then  passing  the  steam  through  a  tube  or  system  of  tubes 
surrounded  by  cold  water.  Thus  the  steam  is  con- 
densed, and  the  distilled  water  is  approximately  pure.  Of 
course,  it  is  necessary  that  the  tubes  in  which  the  con- 
densation takes  place  should  be  of  such  material  that 
water  does  not  act  upon  it  to  any  extent.  The  materials 
used  are  glass,  block  tin,  and  platinum.  Chemically  pure 
water  is  a  very  rare  substance  even  in  the  best  chemical 
laboratories.  The  slight  impurities  present  in  ordinary 
distilled  water  are  not,  however,  of  special  importance 
under  ordinary  circumstances. 


CHAPTER  VI. 

CONSTITUTION    OF   MATTER-ATOMIC    THEORY-ATOMS 
AND  MOLECULES— CONSTITUTION— VALENCE. 

Early  Views. — In  early  times  two  views  were  held  re- 
garding the  ultimate  constitution  of  matter.  The  first 
was,  that  matter  is  infinitely  divisible — that  there  is  no 
limit  to  the  process  of  subdivision ;  the  other  was,  that 
there  is  a  limit  to  the  divisibility — that  when  certain  in- 
conceivably small  particles  are  reached  the  process  must 
stop.  These  small  particles  were  called  atoms,  meaning 
indivisible.  As  long  as  the  constitution  of  matter  was 
merely  a  subject  of  speculation,  the  atoms  remained  with- 
out a  physical  basis  and  were  only  metaphysical  con- 
ceptions. The  facts  which  come  under  our  ordinary  ob- 
servation do  not  furnish  any  evidence  for  or  against  the 
existence  of  atoms ;  and,  though  we  discuss  the  subject 
indefinitely,  little  or  no  progress  can  be  made  without  re- 
fined observations  on  the  properties  of  matter.  So  it  was, 
that  until  the  beginning  of  the  present  century  the  atomic 
theory  remained  practically  what  it  was  when  first  pro- 
posed, and  as  such  it  was  of  no  value  to  chemistry. 

The  Atomic  Theory  as  proposed  by  Dalton. — We  have 
seen  how  Dalton,  at  the  beginning  of  this  century,  dis- 
covered the  law  of  multiple  proportions.  This  law,  as 
well  as  that  of  definite  proportions,  required  explanation. 
The  questions  to  be  answered  are :  (1)  Why  do  the  ele- 
ments combine  in  definite  proportions  ?  (2)  Why,  when 
elements  combine  with  each  other  in  more  than  one  way, 
do  the  relative  quantities  which  enter  into  combination 
in  the  different  cases  bear  simple  relations  to  one  another  ? 
and  (3)  What  is  the  significance  of  the  figures  represent- 
ing the  combining  weights  ?  Dalton  saw  that  the  facts  re- 
ferred to  could  be  explained  by  the  atomic  theory,  while, 
on  the  theory  that  matter  is  infinitely  divisible,  they 

(68) 


THE  ATOMIC  THEORY.  69 

appear  to  be  inexplicable.  It  is  only  necessary  to  assume 
that  each  element  is  made  up  of  particles  which  are  not 
divisible  in  chemical  processes,  and  that  these  particles, 
or  atoms,  have  definite  weights.  The  atoms  of  any  one 
element  must  be  supposed  to  have  the  same  weight, 
while  the  atoms  of  different  elements  have  different 
weights.  Now,  when  chemical  combination  takes  place, 
Dalton  supposed  that  the  action  was  between  the  atoms. 
The  simplest  case  is  that  in  which  combination  takes  place 
in  such  way  that  each  atom  of  one  element  combines 
with  one  atom  of  another ;  but,  besides  this  kind  of  com- 
bination, we  may  have  that  in  which  one  atom  of  one  ele- 
ment combines  with  two  atoms  of  another,  or  two  of  one 
may  combine  with  three  of  another,  etc.  Suppose  two 
elements  A  and  B,  the  weights  of  whose  atoms  are  to  each 
other  as  1 : 10,  are  brought  together,  and  they  combine  in 
the  simplest  way,  i.e.,  one  atom  of  one  with  one  atom  of 
the  other,  then  it  is  plain  that  in  the  compound  AB  the 
elements  will  be  contained  in  the  proportion  of  1  part  of 
A  to  10  parts  of  B,  whether  a  small  or  a  large  quantity 
of  the  compound  is  formed,  and  no  matter  in  what  pro- 
portions the  elements  are  brought  together.  If  they 
should  be  brought  together  in  the  proportion  of  their 
atomic  weights  (1 : 10),  then  no  part  of  either  element  will 
be  left  uncombined  after  the  act  of  combination  has  taken 
place.  If,  however,  a  larger  proportion  of  either  element 
is  taken  than  that  stated,  then  the  quantity  of  the  one 
which  is  in  excess  of  this  proportion  will  be  left  uncom- 
bined. This  is  in  accordance  with  what  we  know  takes 
place,  and  it  is  a  conclusion  drawn  from  the  theory.  No 
matter  how  many  atoms  of  A  we  may  take,  the  same 
number  of  atoms  of  B  will  be  required  to  combine  with 
all  of  them.  But  each  atom  of  B  weighs  10  times  as 
much  as  each  atom  of  A,  therefore  the  total  mass  of  B 
which  enters  into  combination  must  be  10  times  that  of 
A  with  which  it  combines.  It  may  be,  however,  that 
these  same  elements  can  form  other  compounds  with  each 
other.  If  A  and  B  represent  the  atoms  of  the  elements 
and  we  assume  these  atoms  to  be  chemically  indivisible, 
then  the  other  compounds  must  be  represented  by  such 


70  INORGANIC  CHEMISTRY. 

symbols  as  AB»  A^B,  AB^  A^B,  A^B^  etc.,  which  repre- 
sent compounds  in  which  1  atom  of  A  is  combined  with 

2  atoms  of  B ;  2  of  A  with  1  of  B ;  1  of  A  with  3  of  B ; 

3  of  A  with  1  of  B ;  2  of  A  with  3  of  B ;  etc. :  or  they 
also  represent  compounds  in  wrhich  1  part  by  weight  of 
A  is  combined  with  20  parts  by  weight  of  B ;    2  parts  of 
A  with  10  of  B ;    1  of  A  with  30  of  B ;   3  of  A  with  10  of 
B  ;   2  of  A  with  30  of  B ;   etc.     It  is  therefore  clear  that, 
if  the  atomic  theory  as  put  forward  by  Dalton  is  true,  the 
elements  must  combine  according  to  the  laws  of  definite 
and  multiple  proportions  ;  and  it  appears  that  the  figures 
which  represent  the  combining  weights  of  the  elements 
must  either  bear  to  one  another  the  same  relation  as  the 
weights  of  the  atoms,  or,  at  all  events,  the  atomic  weights 
or  the  relative  weights  of  the  atoms  must  be  closely  re- 
lated to  the  combining  weights,  as  will  be  pointed  out 
more  clearly  presently. 

Use  and  Value  of  a  Theory. — The  relation  of  a  theory 
to  facts  is  very  simple,  but  is  frequently  misunderstood. 
The  relation  may  be  conveniently  illustrated  by  the  case 
under  consideration.  By  a  careful  investigation  of  a 
number  of  chemical  compounds  it  was  shown  that  in  each 
of  them  the  same  elements  always  occurred  in  the  same 
proportion.  This  led  to  the  belief  that  this  is  true  of 
every  chemical  compound,  and  after  further  investigation 
which,  as  far  as  it  went,  showed  the  surmise  to  be  correct, 
the  law  of  definite  proportions  was  proposed.  This  law 
is  simply  a  statement  of  what  has  been  found  true  in  all 
cases  examined.  It  involves  no  speculation.  It  is  a 
statement  of  fact.  It  may  be  said  that  the  statement  or 
law  must  be  open  to  some  doubt  for  the  reason  that  all 
possible  cases  have  not  been  examined,  and  it  may  not 
hold  true  for  some  of  these  unexamined  cases.  The  reply 
to  this  is  that  it  has  been  found  true  in  a  very  large  num- 
ber of  cases  and  in  all  cases  which  have  been  investigated. 
It  is  true  for  the  present  state  of  our  knowledge,  and  that 
is  all  we  can  demand  of  any  law.  Again,  further  investi- 
gation led  to  the  discovery  of  the  law  of  multiple  propor- 
tions, which  is  also  a  statement  of  what  has  been  found 
true  in  all  cases  investigated.  It,  like  the  law  of  definite 


ATOMIC  WEIGHTS  AND   COMBINING  WEIGHTS.        71 

proportions  and  in  the  same  sense,  is  a  statement  of  fact. 
But  having  gone  thus  far,  we  now  ask,  what  is  the  ex- 
planation of  these  laws?  We  simply  know  the  facts — 
what  is  the  explanation  ?  By  experiment  we  cannot  go 
beyond  these  facts,  but  it  is  possible  to  imagine  a  cause 
and  then  proceed  to  see  whether  the  imagined  cause  is 
sufficient  to  account  for  the  facts.  This  is  what  Dalton 
did.  He  imagined  that  matter  is  made  up  of  atoms  of 
definite  weights,  and  that  chemical  combination  takes 
place  in  simple  ways  between  these  atoms.  This  imag- 
ined cause  is  the  atomic  theory.  It  is  not  a  statement  of 
anything  found  by  investigation.  It  is  not  a  statement 
of  an  established  fact.  It  may  or  may  not  be  literally 
true,  but  at  all  events  it  is  the  best  guess  that  has  ever 
been  made  as  to  the  cause  of  the  fundamental  laws  of 
chemical  action,  and  it  furnishes  a  very  convenient  means 
of  interpreting  the  facts  of  chemistry.  Since  the  atomic 
theory  was  first  proposed  it  has  been  accepted  by  nearly 
all  chemists.  It  has  been  of  great  value  in  suggesting 
methods  of  work,  and  has  contributed  largely  to  the  ad- 
vance of  chemistry.  Any  theory  which  is  in  accord- 
ance with  the  facts  and  leads  to  the  discovery  of  new 
facts  is  of  value,  whether  it  should  eventually  prove 
to  be  true  or  false.  At  the  same  time  a  false  theory 
may  do  much  harm,  as  it  may  lead  men  to  misinterpret 
the  facts  which  they  observe,  and  thus  retard  progress. 

Atomic  Weights  and  Combining  "Weights. — If  the  atomic 
theory  is  true  the  atoms  of  each  element  must  have  defi- 
nite weights,  and  the  determination  of  these  atomic 
weights  must  evidently  be  of  great  importance.  By 
analysis  of  compounds  we  can  only  determine  the  pro- 
portions by  weight  in  which  the  elements  combine  with 
one  another.  Can  we  in  this  way  determine  the  atomic 
weights  ?  In  the  first  place,  it  is  clear  that  it  is  out  of 
the  question  to  think  of  determining  the  absolute  weights 
of  the  atoms,  and  all  that  we  can  possibly  do  is  to  deter- 
mine their  relative  weights.  As  of  all  the  elements  hy- 
drogen enters  into  combination  in  smallest  relative  quan- 
tity, its  atomic  weight  is  taken  as  the  unit  of  the  system, 
and  the  problem  before  us  is  to  determine  how  many 


72  INORGANIC  CHEMISTRY. 

times  heavier  the  atoms  of  the  other  elements  are  than 
that  of  hydrogen.  If  every  element  combined  with  hydro- 
gen in  only  one  proportion  the  problem  would  be  a  com- 
paratively simple  one.  Thus  the  three  elements  chlorine, 
bromine,  and  iodine  combine  with  hydrogen,  forming  only 
one  compound  each.  On  analysis  these  are  found  to 
contain  respectively 

1  part  of  hydrogen  to    35.18  parts  of  chlorine  ; 

1     "  "          "     79,34      "      "  bromine;  and 

1     "  "          "  125.89      "      "  iodine. 

There  is  no  reason  for  believing  that  in  these  compounds 
the  elements  are  combined  in  any  but  the  simplest  way, 
i.e.,  that  each  atom  of  hydrogen  is  combined  with  one 
atom  of  chlorine  to  form  hydrochloric  acid,  etc.  If  this 
is  true,  then  the  atom  of  chlorine  must  weigh  35.37  times  ; 
that  of  bromine  79.76  times;  and  that  of  iodine  126.54 
times  as  much  as  that  of  hydrogen,  or,  in  other  words, 
the  atomic  weights  of  chlorine,  bromine,  and  iodine  are 
respectively  35.37,  79.76,  and  126.54.  It  will,  however, 
be  observed  that  there  is  no  evidence  as  to  whether  the 
elements  in  these  compounds  are  combined  in  the  simplest 
way  or  not.  It  is  possible,  as  far  as  we  know,  that  one 
atom  of  hydrogen  may  combine  with  two  or  three  of 
chlorine,  or  that  one  of  chlorine  may  combine  with  two 
or  three  of  hydrogen.  As  there  is,  however,  no  evidence 
upon  this  point  the  simplest  assumption  is  made. 

If  we  take  the  case  of  oxygen  the  problem  is  more  com- 
plex. In  water  the  elements  are  combined  in  the  propor- 
tion of  7.94  parts  of  oxygen  to  1  part  of  hydrogen,  and 
from  this  we  should  naturally  conclude  that  the  atomic 
weight  of  oxygen  is  7.94 ;  but  further  study  shows  that 
this  conclusion  is  not  justified.  Hydrogen  and  oxygen 
form  a  second  compound  known  as  hydrogen  dioxide  or 
hydrogen  peroxide  in  which  there  are  15.88  parts  of  oxy- 
gen to  1  of  hydrogen.  This  may  be  explained  in  the 
terms  of  the  atomic  theory  by  assuming  that  water  is 
represented  by  the  formula  HO,  and  hydrogen  dioxide 
by  HO3.  But  it  may  be  that  the  atomic  weight  of  oxygen 
is  15.88,  and  then  water  must  be  represented  by  the 


MOLECULES- AVOGADRO'S  LAW.  73 

formula  H2O,  and  hydrogen  dioxide  by  HO.  Simple 
analysis  of  the  compounds  is  not  sufficient  to  enable  us  to 
decide  between  these  possibilities.  It  is,  therefore,  evi- 
dent that  in  order  to  determine  the  atomic  weights  some- 
thing besides  the  determination  of  the  composition  of 
compounds  is  necessary.  The  figures  representing  the 
combining  weights  found  in  this  way  will,  however,  either 
be  identical  with  the  atomic  weights  or  will  bear  a  simple 
numerical  relation  to  them. 

Molecules. — Investigation  of  certain  phenomena  of 
light,  of  electricity,  of  liquid  films  and  the  conduct  of 
gases  has  led  physicists  to  the  conclusion  that  matter  is 
not  continuous,  but  made  up  of  small  particles,  which  are 
called  molecules.  A  gaseous  molecule  is  defined  as  "  that 
minute  portion  of  a  substance  which  moves  about  as  a  whole, 
so  that  its  parts,  if  it  has  any,  do  not  part  company  during 
the  motion  of  agitation  of  the  gas."  It  would  be  out  of 
place  here  to  present  the  physical  facts  upon  which  the 
molecular  theory  rests.  Suffice  it  to  say  that  it  is  the 
only  theory  which  has  been  found  adequate  to  account 
for  the  behavior  of  gases. 

Avogadro's  Law. — The  fact  that  gases  conduct  them- 
selves in  the  same  way  under  the  influence  of  changes  in 
temperature  and  pressure  can  only  be  explained  by  as- 
suming that  equal  volumes  of  all  gases  and  vapors  contain 
the  same  number  of  ultimate  particles  or  molecules  at  the 
mme  temperature  and  pressure. 

This  is  a  deduction  from  the  well-tested  molecular 
theory  of  gases.  It  was,  however,  originally  put  forward 
by  the  Italian  chemist  Avogadro  from  a  study  of  chemical 
as  well  as  of  physical  facts,  and  a  little  later  it  was  sug- 
gested as  probable  by  the  French  physicist  Ampere.  It 
is  therefore  generally  spoken  of  as  Avogadro's  law,  and 
sometimes  as  Ampere's  law.  Absolute  proof  of  its  truth 
cannot  be  given,  but  it  is  in  thorough  accordance  with  a 
la*rge  number  of  wTell-known  facts,  and  it  is  undoubtedly 
true,  if  the  molecular  theory  of  matter  is  true.  It  may 
therefore  be  considered  as  furnishing  a  solid  foundation 
for  further  conclusions  bearing  upon  the  problem  of  the 
•determination  of  the  atomic  weights. 


74  INORGANIC  CHEMISTRY. 

Distinction  between  Molecules  and  Atoms. — If  we  con- 
sider any  chemical  compound,  as  water  or  hydrochloric 
acid,  it  is  evident  that  the  smallest  particle  or  the  mole- 
cule of  the  compound  must  be  made  up  of  still  smaller 
particles.  Thus,  the  smallest  particle  of  water  must  con- 
tain smaller  particles  of  hydrogen  and  oxygen,  and  the 
smallest  particle  of  hydrochloric  acid  must  contain  smaller 
particles  of  hydrogen  and  chlorine.  These  smallest  par- 
ticles of  the  molecules  are  the  atoms.  The  molecules  of 
the  compounds  are,  according  to  this  view,  made  up  of 
the  atoms  of  the  elements.  Similarly  the  elements  them- 
selves are,  for  good  reasons  which  will  be  presented,  be- 
lieved to  consist  of  molecules  which  are  in  turn  made  up 
of  -atoms  of  the  same  kind,  though  in  a  few  cases  the 
molecule  of  the  element  is  identical  with  the  atom.  The 
difference  between  a  compound  and  an  element  then  is, 
in  general,  that  the  molecule  of  the  compound  consists  of 
atoms  of  different  kinds,  while  the  molecule  of  an  ele- 
ment consists  of  atoms  of  the  same  kind  or,  in  a  few 
oases,  of  one  atom.  Generally  the  atoms  do  not  exist  in 
the  free  or  uncombined  state,  but,  if  they  are  set  free  by 
chemical  action,  they  unite  to  form  molecules.  The  fol- 
lowing may  serve  as  a  definition  of  the  conception  of 
atoms  at  present  held  by  chemists : 

Atoms  are  the  indivisible  constituents  of  molecules.  They 
are  the  smallest  particles  of  the  elements  that  take  part  in 
chemical  reactions,  and  are,  for  the  greater  part,  incapable 
of  existence  in  the  free  state,  being  generally  found  in  combi- 
nation with  other  atoms,  either  of  the  same  kind  or  of  differ- 
ent kinds. 

It  cannot  be  too  strongly  emphasized  that  the  views 
held  in  regard  to  the  relations  between  molecules  and 
atoms  are  based  upon  an  enormous  amount  of  painstak- 
ing study  of  facts,  and  in  order  fully  to  comprehend  their 
value  a  study  of  most  of  these  facts  would  be  necessary. 
These  views  have  gradually  become  firmly  established  as 
knowledge  of  the  facts  has  grown  more  and  more  pro- 
found. Accepting  them,  we  are  now  to  see  how  they  aid 
us  in  the  problem  with  which  we  are  dealing,  viz.,  the 
determination  of  the  atomic  weights. 


MOLECULAR  WEIGHTS.  75 

Molecular  Weights. — If  equal  volumes  of  gases  contain 
the  same  number  of  molecules  at  the  same  temperature 
and  pressure,  it  is  only  necessary  to  determine  the 
weights  of  equal  volumes  of  gases  to  learn  the  relative 
weights  of  their  molecules.  Thus,  if  we  weigh  a  liter  of 
each  of  three  gases,  and  find  that  the  weights  are  to  one 
another  as  1  to  2  to  3,  then  it  follows  that  the  relation 
between  the  weights  of  the  molecules  of  these  gases  is 
expressed  by  these  figures,  or,  in  other  words,  the  mole- 
cule of  the  second  gas  is  twice  as  heavy ;  and  that  of  the 
third  gas  is  three  times  as  heavy  as  that  of  the  lightest. 
The  determination  of  the  relative  weights  of  the  mole- 
cules of  substances  which  either  are  gaseous  or  can  be 
converted  into  gases  resolves  itself  simply  into  a  deter- 
mination of  the  weights  of  equal  volumes.  In  represent- 
ing the  molecular  weights  we  may  use  any  figures  which 
are  most  convenient,  provided  only  that  they  bear  to  one 
another  the  relations  determined  by  experiment.  If, 
however,  we  call  the  atomic  weight  of  hydrogen  1,  then 
our  system  of  molecular  weights  must  be  based  upon 
this,  and  the  molecular  weight  of  a  compound  should 
state  how  much  heavier  the  molecule  is  than  an  atom  of 
hydrogen.  Thus,  if  we  say  that  the  molecular  weights 
of  water  and  hydrochloric  acid  are  respectively  17.88  and 
36.18,  we  mean  that,  if  the  hydrogen  atom  weighs  1,  then 
the  weights  of  the  molecules  of  water  and  hydrochloric 
acid  are  represented  by  the  figures  given.  Now,  if  the 
molecule  of  hydrogen  were  identical  with  the  atom  or,  in 
other  words,  if  the  molecular  weight  were  equal  to  1,  then 
the  adjustment  of  the  system  of  molecular  weights  to  the 
atomic  weight  of  hydrogen  would  be  perfectly  simple. 
It  would  only  be  necessary  to  determine  the  weight  of  a 
given  volume  of  hydrogen  and  compare  the  weights  of 
equal  volumes  of  other  gases  with  it.  If  the  weight  of  a 
certain  volume  of  any  compound  should  be  found  to  be 
10  times  that  of  an  equal  volume  of  hydrogen,  it  would 
follow  that  the  molecular  weight  of  the  compound  is  10. 
But  the  molecule  and  atom  of  hydrogen  are  not  identical, 
as  can  be  shown  without  difficulty.  When  a  given  vol- 
ume of  hydrogen  combines  with  chlorine  it  combines  with 


76 


INORGANIC  CHEMISTRY. 


an  equal  volume  of  this  element,  and  the  two  volumes 
which  combine  form  an  equal  volume  of  the  compound 
hydrochloric  acid.  These  facts  may  be  graphically  repre- 
sented as  follows  : 


combine  and  form 


1vol. 

HC1 

2   volumes   of 

-    hydrochloric 
acid  gas. 

1vol. 

HC1 

1 

j 

Now,  bearing  in  mind  Avogadro's  law  that  equal  vol- 
umes of  all  gases  contain  the  same  number  of  molecules, 
it  follows  that  if,  in  the  volume  of  hydrogen  taken,  there 
is  any  finite  number  of  molecules,  say  100,  then  in  the 
same  volume  of  chlorine  there  must  be  100  molecules, 
and  in  the  two  volumes  of  hydrochloric  acid  gas  obtained 
there  must  be  200  molecules  of  the  compound.  There- 
fore, from  100  molecules  of  hydrogen  and  100  molecules 
of  chlorine  there  are  formed  200  molecules  of  hydro- 
chloric acid.  But  in  each  molecule  of  hydrochloric  acid 
there  must  be  at  least  one  atom  of  hydrogen  and  one 
atom  of  chlorine,  and  in  the  200  molecules  there  must  be 
at  least  200  atoms  of  hydrogen  and  200  atoms  of  chlo- 
rine. Now,  these  200  atoms  of  hydrogen  have  come  from 
the  100  molecules,  and  the  same  is  true  of  the  chlorine. 
It,  therefore,  follows  that  each  molecule  of  hydrogen  and 
each  molecule  of  chlorine  must  consist  of  at  least  two  atoms. 
Or,  we  may  say  that,  if  there  is  one  atom  of  hydro- 
gen in  the  molecule  of  hydrochloric  acid,  then  there 
are  two  atoms  of  hydrogen  in  the  molecule  of  hydrogen. 
The  assumption  that  the  molecule  of  hydrogen  is  twice 
as  heavy  as  the  atom  is  found  to  satisfy  every  require- 
ment of  the  present  state  of  our  knowledge.  From  the 
above,  the  following  rule  for  the  determination  of  molec- 
ular weights  is  deduced  :  Determine  the  specific  gravity 
of  the  substance  in  terms  of  hydrogen  and  multiply  the 


DEDUCTION  OF  ATOMIC  WEIGHTS. 


result  by  ,2.  Tims  the  specific  gravity  of  water  vapor  in 
terms  of  hydrogen  is  8.94,  or,  in  other  words,  a  given  vol- 
ume of  water  vapor  weighs  8.94  times  as  much  as  the 
same  volume  of  hydrogen.  But  the  molecular  weight  of 
hydrogen  being  2,  that  of  water  must  be  17.88.  As  the 
specific  gravity  of  gases  is  frequently  stated  in  terms  of 
the  air  standard,  it  is  desirable  to  know  the  relation  be- 
tween these  figures  and  those  based  upon  hydrogen.  The 
specific  gravity  of  hydrogen  as  compared  with  air  is 
€.06926 ;  when  taken  as  the  standard  it  is  represented  by 
1.  But  1  -f-  0.06926  is  14.44 ;  therefore,  to  convert  the 
specific  gravities  on  the  air  standard  into  those  on  the 
hydrogen  standard  it  is  only  necessary  to  multiply  by 
14.44.  Thus  the  specific  gravity  of  water  vapor,  air  be- 
ing the  standard,  is  0.623.  To  find  its  specific  gravity,  hy- 
drogen being  the  standard,  multiply  by  14.44.  Now  0.623 
X  14.44  is  very  nearly  9.  This  being  the  specific  gravity, 
the  molecular  weight  is  obtained  by  multiplying  by  2. 
Of  course,  we  should  reach  the  same  result  by  multiply- 
ing the  specific  gravities  in  terms  of  air  directly  by  28.88. 
Below  are  given  the  molecular  weights  of  a  few  elements 
and  compounds  which  have  been  determined  by  the 
method  just  described : 


Name. 

Sp.  Grav. 
H  =  1. 

Molec. 
Weight. 

Molecular 
Formula. 

Hydrogen  

1 

2 

Ho 

Nitrogen  

13  93 

27  86 

N2 

Water  „  

8  94 

17  88 

H2O 

Hydrochloric  acid  

18  09 

36  18 

Hen 

Ammonia  ,  

8  46 

16  93 

NH3 

7  96 

15  92 

CHd 

Carbon  monoxide  „  ,  

13  90 

27  80 

CO 

Carbon  dioxide  

21  84 

43  68 

CO3 

Deduction  of  Atomic  Weights  from  Molecular  Weights. — 
The  determination  of  molecular  weights  does  not  neces- 
sarily carry  with  it  the  determination  of  the  atomic 
weights.  It  is  plain  from  what  has  already  been  said 
that  a  knowledge  of  the  molecular  weight  of  an  element 
does  not  convey  a  knowledge  of  its  atomic  weight.  If, 


78  INORGANIC  CHEMISTRY. 

for  example,  we  learn  that  the  molecular  weight  of  nitro- 
gen is  approximately  28,  we  have  no  means  of  judging 
from  this  what  the  atomic  weight  is.  It  is  plainly  nec- 
essary to  know  of  how  many  atoms  each  molecule  of 
nitrogen  is  made  up,  and  to  learn  this  is  not  a  simple 
matter.  It  is  easier  to  determine  the  atomic  weight  of 
an  element  through  a  study  of  its  compounds.  Suppose 
it  is  desired  to  determine  the  atomic  weight  of  oxygen. 
We  first  determine  the  molecular  weights  of  a  number  of 
compounds  which  contain  oxygen,  and  then  analyze  these 
compounds.  We  then  see  what  the  smallest  figure  is 
that  is  required  to  express  the  weight  of  the  oxygen 
that  enters  into  the  composition  of  the  molecules,  and 
this  figure  is  selected  as  the  atomic  weight.  The  mo- 
lecular weights  and  the  composition  of  several  oxygen 
compounds  are  given  in  the  following  table  : 

Compound.                                       Mol.  Wt.  Approx.  Composition. 

Water 17.88  2       parts  hydrogen, 


15.88 


Carboii  monoxide 27.80  11.92  carbon, 


15.88 


Carbon  dioxide 43.68  11.92  carbon, 


31.76 


Nitric  oxide 29.81  13.93  nitrogen 


15.88 


Nitrous  oxide 43.74  27.86  nitrogen, 


15.88 


Sulphur  dioxide 63.59  31.83  sulphur, 


31.76 


Sulphur  trioxide 79.47  31.83  sulphur, 


47.64 


oxygen. 


oxygen. 


oxygen. 


oxygen. 


oxygen. 


oxygen. 


oxygen. 


The  figures  in  the  third  column  are  of  course  determined 
by  analysis,  an  example  of  the  methods  used  having  been 
given  in  the  chapter  on  water.  Stated  in  ordinary  lan- 
guage, the  figures  in  the  case  of  carbon  monoxide  mean  that 
the  molecule  of  this  compound  weighs  27.80  times  as  much 
as  the  atom  of  hydrogen,  and  the  27.80  parts  of  matter 
are  made  up  of  11.92  parts  of  carbon  and  15.88  parts  of 
oxygen.  Considering  now  the  composition  of  the  com- 
pounds in  the  table,  it  will  be  seen  that  the  smallest  mass 
of  oxygen  which  enters  into  the  composition  of  any  of  the 


'MOLECULAR  FORMULAS.  79 

molecules  weighs  15.88  times  as  much  as  the  atom  of 
hydrogen.  We  find  twice  this  mass  as  in  carbon  dioxide 
and  sulphur  dioxide  ;  and  three  times  as  in  sulphur  tri- 
oxide,  but  no  smaller  mass.  Now,  if  we  should  examine 
all  compounds  of  oxygen  which  can  exist  in  the  form  of 
gas  or  vapor  we  should  find  the  same  thing  true  ;  that  is 
to  say,  the  smallest  mass  of  oxygen  which  enters  into  the 
composition  of  molecules  is  15.88  as  great  as  that  of  the 
atom  of  hydrogen.  The  conclusion  is  therefore  drawn 
that  15.88  is  the  atomic  weight  of  oxygen.  The  possi- 
bility that  the  atomic  weight  of  oxygen  is  less  than  this 
figure  is  not  excluded.  It  may  be  that  in  the  simplest 
oxygen  compounds  now  known  there  are  two  or  more 
atoms  of  this  element  in  the  molecules.  But  in  the  total 
absence  of  evidence  on  this  point  all  we  can  do  is  to 
accept  the  figure  15.88  as  in  perfect  accordance  with  all 
our  knowledge  of  oxygen  compounds. 

In  this  way  the  atomic  weights  of  all  elements  which 
form  gaseous  compounds  or  compounds  that  can  be  con- 
verted into  vapor  have  been  determined  ;  and  the  deter- 
minations made  in  this  way  are  regarded  as  the  most 
reliable. 

Exact  Atomic  Weights  determined  by  the  Aid  of  Analy- 
sis.— By  determining  molecular  weights  it  is  possible 
to  decide  approximately  what  figure  represents  the  atomic 
weight  of  an  element,  but  the  methods  employed  in  making 
determinations  of  molecular  weights  are  liable  to  slight 
errors,  and  therefore  the  atomic  weights  obtained  directly 
from  the  molecular  weights  deviate  slightly  from  the  true 
figures.  In  order  to  determine  the  atomic  weights  with 
the  greatest  possible  accuracy,  the  most  refined  methods 
of  chemical  analysis  are  brought  into  play,  and  the  figures 
in  the  table  on  page  21  have  been  determined  in  this  way 
by  a  combination  of  a  study  of  the  specific  gravity  of 
gases  and  by  the  most  careful  analyses,  together  with 
some  other  methods  which  will  be  taken  up  later. 

Molecular  Formulas. — The  symbols  of  chemical  com- 
pounds first  used  were  intended  to  express  simply  the 
composition  of  the  compounds,  and  this  can  be  done  as 
was  explained  in  Chapter  I.  by  adopting  a  system  of 


80  INORGANIC  CHEMISTRY. 

combining  weights  of  the  elements.  According  to  the 
theory  explained  in  the  last  chapter  the  smallest  particle 
of  every  compound  is  a  molecule,  and  each  molecule  is 
made  up  of  atoms.  It  appears,  therefore,  desirable  for 
the  sake  of  uniformity  that  the  symbols  used  to  repre- 
sent chemical  compounds  should  represent  molecules. 
Where  the  molecular  weight  of  a  compound,  the  atomic 
weights  of  the  elements  of  which  it  is  composed,  and  its 
composition  are  known,  there  is  no  difficulty  in  represent- 
ing it  by  a  molecular  formula.  Thus,  the  molecular 
weight  of  ammonia  is  found  by  experiment  to  be  approxi- 
mately 17,  and  the  17  parts  are  made  up  of  14  parts  of 
nitrogen  and  3  parts  of  hydrogen.  The  atomic  weight 
of  nitrogen  is  found  by  the  method  which  has  just  been 
described  to  be  very  nearly  14.  Therefore  the  mole- 
cule of  ammonia  weighing  17  parts  is  composed  of  1 
atom  of  nitrogen  weighing  14  parts  and  3  atoms  of  hy- 
drogen weighing  3  parts.  The  composition  of  the  mole- 
cule is  therefore  represented  by  the*  formula  NH3.  Simi- 
larly the  composition  of  the  molecule  of  water  is  repre- 
sented by  the  formula  H2O ;  that  of  hydrochloric  acid 
by  HOI ;  that  of  marsh  gas  by  CH4 ;  etc.,  etc.  Every 
formula  now  in  use  is  intended  to  represent  a  molecule  of 
the  compound  for  which  it  stands.  In  regard  to  the 
molecular  weights  of  compounds  that  are  not  gaseous 
nor  convertible  into  vapor,  Avogadro's  method  is  plainly 
of  no  avail.  Methods  have,  however,  been  devised  which 
are  applicable  to  a  number  of  these  (see  Chapter  XXIII). 
Constitution. — When  hydrochloric  acid  is  formed,  we 
conceive  that  each  atom  of  hydrogen  combines  with  one 
atom  of  chlorine,  and  that  the  molecules  of  the  resulting 
compound  are  made  up  each  of  an  atom  of  hydrogen  and 
an  atom  of  chlorine.  What  the  act  of  combination  con- 
sists in  we  do  not  know.  We  simply  know  that  something 
very  remarkable  takes  place,  and  that  as  a  consequence 
the  hydrogen  and  chlorine  cease  to  exist  in  their  original 
forms.  It  is  idle  at  present  even  to  speculate  in  regard 
to  the  character  of  the  change.  The  fact  of  union  is  ex- 
pressed by  writing  the  symbols  of  the  elements  side  by 
side  without  any  sign  between  them,  as  HC1,  or;  some- 


VALENCE.  81 

times,  it  is  convenient  to  use  a  line  to  indicate  chemical 
union,  thus :  H-C1.  According  to  the  molecular  theory 
the  molecule  of  water  consists  of  two  atoms  of  hydrogen 
and  one  of  oxygen,  as  represented  by  the  formula  H2O, 
and  the  question  now  suggests  itself  whether  all  three 
atoms  are  in  combination  with  one  another  or  whether 
each  of  the  hydrogen  atoms  is  in  combination  with  the 
oxygen  atom,  but  not  with  each  other,  as  represented  by 
the  formula  H-O-H.  So  too  in  the  case  of  ammonia, 
the  molecular  formula  of  which  is  NH3,  the  question  sug- 
gests itself :  Are  the  three  atoms  of  hydrogen  in  combi- 
nation with  the  atom  of  nitrogen,  but  not  with  one  an- 

/H 

other,  as  represented  in  the  formula  N^-H  ?     It  is  ex- 
Mi 

tremely  difficult  to  answer  such  questions,  but,  at  the 
same  time,  certain  facts  are  known  which  enable  us  to 
draw  probable  conclusions.  Formulas  which  express  the 
composition  of  molecules  and  at  the  same  time  express 
the  relations  or  the  connections  which  exist  between  the 
atoms  are  called  constitutional  formulas.  These  constitu- 
tional formulas  are  very  frequently  used  at  present,  but 
sometimes  without  a  sufficient  basis  of  facts  to  justify 
them.  Whenever  they  are  used  in  this  book,  the  rea- 
sons for  them  will  be  stated  as  fully  as  may  appear  nec- 
essary. 

Valence. — The  formulas  of  the  hydrogen  compounds 
of  chlorine,  oxygen,  nitrogen,  and  carbon,  all  determined 
by  the  same  method,  are 

C1H        OH,        NH3        CH4. 

A  consideration  of  these  formulas  and  of  many  similar 
ones  has  led  to  the  belief  that  the  atoms  of  different  ele- 
ments differ  in  their  power  of  holding  other  atoms  in 
combination.  The  simplest  explanation  of  the  composi- 
tion of  the  compounds  above  represented  is  that  the 
atoms  of  chlorine,  oxygen,  nitrogen,  and  carbon  differ  in 
their  power  of  holding  hydrogen  atoms  in  combination. 
Hydrogen  and  chlorine  combine  in  only  one  way,  1  atom 
of  chlorine  combining  with  1  of  hydrogen ;  1  of  oxygen 


82  INORGANIC  CHEMISTRY. 

combines  with  2  of  hydrogen;  1  of  nitrogen  with  3  of 
hydrogen ;  and  1  of  carbon  with  4  of  hydrogen.  The 
limit  of  the  combining  power  of  the  atom  of  chlorine  is 
reached  when  it  has  combined  with  one  atom  of  hydro- 
gen. And  as  one  chlorine  atom  can  hold  but  one  atom 
of  hydrogen  in  combination,  so  one  atom  of  hydrogen 
can  hold  but  one  atom  of  chlorine.  Either  the  hydrogen 
atom  or  the  chlorine  atom  may  be  taken  as  an  example 
of  the  simplest  kind  of  atom.  Any  element  like  hydro- 
gen or  chlorine  is  called  a  univalent  element ;  an  element 
like  oxygen  whose  atom  can  hold  two  unit  atoms  in 
combination  is  called  a  bivalent  element ;  an  element  like 
nitrogen  whose  atom  can  hold  three  unit  atoms  in  com- 
bination is  called  a  trivalent  element ;  and  an  element  like 
carbon  whose  atom  can  hold  four  unit  atoms  in  combina- 
tion is  called  a  quadrivalent  element.  Most  elements  be- 
long to  one  or  the  other  of  these  four  classes,  though 
there  are  some  which  can  hold  five,  six,  and  even  seven 
unit  atoms  in  combination.  These  are  called  quinqui- 
valent, sexivalent,  and  septivalent  respectively. 

Valence  is  defined  as  that  property  of  an  element  by 
virtue  of  which  its  atom  can  hold  a  definite  number  of 
other  atoms  in  combination.  In  the  formation  of  com- 
pounds the  valence  of  the  elements  determines  how  many 
atoms  of  any  element  can  enter  into  combination  with 
any  other.  The  atoms  are  sometimes  spoken  of  as  hav- 
ing bonds  which  are  graphically  represented  by  lines. 
Thus,  a  univalent  element  is  said  to  have  one  bond,  as 
represented  by  H-,  C1-,  etc. ;  a  bivalent  element  is  said 
to  have  two  bonds,  -O-,  -S-,  etc.  ;  .a  trivalent  element 

three,  -N- ;   and  a  quadrivalent  element  four,  -C-.     Of 

course,  this  is  merely  a  symbolical  representation  of  the 
idea  that  each  atom  has  a  definite  power  of  combining 
with  others.  It  is  further  said  that  when  the  atoms  unite 
these  bonds  become  satisfied.  Thus  when  one  atom  of 
hydrogen  unites  with  one  of  chlorine,  the  bond  of  each 
is  regarded  as  uniting  with  the  bond  of  the  other,  and  this 
is  represented  by  the  symbol  H-C1.  So  too,  when  two 
atoms  of  hydrogen  unite  with  one  of  oxygen,  the  com- 


REPLACING  POWER  OF  ELEMENTS.  83 

TT 

pound  is  represented  in  this  way  :  H-O-H  or  O<TJ-     In 

the  union  of  atoms,  further,  to  use  the  figurative  lan- 
guage, two  bonds  may  be  satisfied  by  two  univalent  atoms 
or  by  one  bivalent  atom.  Thus,  in  marsh  gas,  CH4,  the 
four  affinities  or  bonds  of  the  quadrivalent  carbon  atom 
are  regarded  as  being  satisfied  by  the  four  univalent 

H 

hydrogen  atoms.  In  the  compound  H-C=O,  however, 
two  of  the  bonds  are  regarded  as  satisfied  by  two  univa- 
lent hydrogen  atoms,  and  the  other  two  by  one  bivalent 
oxygen  atom  ;  and,  again,  in  the  compound  O=C—  O,  the 
four  bonds  of  the  carbon  atom  are  regarded  as  satisfied 
by  two  bivalent  oxygen  atoms. 

Replacing  Power  of  Elements.  —  As  has  been  seen, 
when  potassium  acts  upon  water  the  action  is  represented 
by  the  equation 


Expressing  this  reaction  by  means  of  formulas  which 
take  the  valence  of  the  elements  into  consideration  we 
have  the  following  : 

K  +  H-O-H  =  K-O-H  +  H. 

According  to  this,  one  atom  of  potassium  is  substituted 
for  one  atom  of  hydrogen  in  water,  and,  in  the  compound 
formed,  the  atom  of  potassium  is  regarded  as  occupying 
the  place  of  the  hydrogen  atom.  The  elements  calcium 
and  barium  are  bivalent  like  oxygen,  as  shown  in  the 
compounds  CaCl2  and  BaCl2,  in  which  the  atom  of  cal- 
cium as  well  as  that  of  barium  holds  two  univalent  atoms 
of  chlorine  in  combination.  When  these  elements  act 
upon  water,  hydrogen  is  liberated  as  with  potassium,  but 
each  atom  replaces  two  atoms  of  hydrogen,  as  represent- 
ed in  the  equations 

Ca  +  2H2O  =  CaHaO3  +  2H  ; 
Ba  +  2H2O  =  BaH2Oa  +  2H. 


84  INORGANIC  CHEMISTRY. 

Or  expressing  the  changes  by  valence  formulas,  we  have 


-o     ,  H-O-H      -o      0-H      „ 
Ba  +  H-O-H  =  Ba<0-H  +  H'* 

A  trivalent  element  acting  in  the  same  way  would  give  a 

/0-H 
compound  of  the  general  formula  M(—  O-H,  in  which  M 

\O-H 
represents  any  trivalent  element. 

So,  too,  in  the  action  of  various  elements  upon  hydro- 
chloric acid,  HC1,  a  univalent  element  like  sodium  or 
potassium  replaces  one  hydrogen  atom  in  one  molecule 
of  the  acid,  and  forms  a  compound  of  the  general  formula 
MCI,  as,  for  example,  KC1  and  NaCl.  A  bivalent  ele- 
ment, like  zinc,  replaces  two  atoms  of  hydrogen  in  two 
molecules  of  the  acid,  forming  the  compound  ZnCl2  : 


A  trivalent  element  forms  a  compound  of  the  general 
formula  MC13,  as,  for  example,  aluminium  chloride  : 

/Cl 

AlCl. 


The  above  explanation  of  the  hypothesis  of  valence  will 
suffice  by  way  of  introduction  to  the  subject.  The  re- 
lation which  it  bears  to  the  facts  is  simply  this  :  On  study- 
ing the  composition  of  chemical  compounds  we  find  that, 
in  general,  the  elements  combine  with  one  another  in 
comparatively  few  proportions.  Thus,  hydrogen  and 
chlorine  combine  with  each  other  in  only  one  proportion  ; 
hydrogen  and  oxygen  in  two  ;  nitrogen  and  hydrogen 
in  four  ;  etc.  There  is  something  limiting  the  com- 
plexity of  compounds.  We  might  study  the  laws  govern- 
ing the  complexity  of  compounds  without  any  hypothesis 
as  to  the  cause,  but  the  hypothesis  of  valence  is  a  con- 
venient explanation  of  these  laws,  and  it  has  been  of  much 
service  in  furnishing  chemists  with  a  simple  language  for 
representing  chemical  changes. 


CHAPTER  VII. 

OZONE  —  ALLOTROPY  —  NASCENT    STATE  —  HYDROGEN 
DIOXIDE. 

Occurrence. — Ozone  lias  long  been  thought  to  be  pres- 
ent in  the  air  in  small  quantity,  but  careful  research  has 
made  this  occurrence  appear  doubtful. 

Preparation. — It  was  observed  in  the  last  century  that, 
when  a  powerful  electric  machine  is  worked  in  a  room, 
something  is  formed  which  has  a  strong  odor,  and  the 
same  odor  was  noticed  during  thunder-showers.  After- 
wards the  same  thing  was  noticed  when  water  is  decom- 
posed by  an  electric  current,  and  when  phosphorus  is  ex- 
posed to  moist  air.  The  substance  was  at  first  supposed 
to  be  a  compound  of  water  and  oxygen,  but  long-contin- 
ued investigation  showed  that  it  could  be  made  from  pure 
dry  oxygen,  and  that  by  heat  and  other  means  it  is  de- 
composed into  nothing  but  ordinary  oxygen. 

It  is  prepared  mixed  with  oxygen  by  passing  electric 
sparks  through  ordinary  oxygen  in  an  apparatus  con- 
structed on  the  principle  of  that  represented  in  Fig.  4. 


FIG.  4. 

AA  is  a  glass  tube  about  an  inch  in  diameter  closed  at 
the  ends  with  brass  caps  or  with  corks,  covered  with  shel- 
lac on  the  inner  side.  A  metallic  cylinder  BB  covered 
with  tin-foil  is  placed  inside  this  tube.  The  cylinder  is 
connected  with  the  tubes  (70,  and  through  these  and  the 

85 


86  INORGANIC  CHEMISTRY. 

cylinder  a  current  of  cold  water  is  kept  flowing.  The 
oxygen  passes  through  the  glass  tube  by  means  of  the 
tubes  DD,  and  necessarily  passes  through  the  narrow 
space  between  the  glass  tube  and  the  metallic  cylinder. 
Around  the  outside  of  the  glass  tube  is  wound  a  strip  of 
tin-foil.  By  means  of  the  wires  F  and  E  connection  is 
established  with  the  poles  of  an  induction-coil,  or  a  Holtz 
electrical  machine. 

Ozone  is  also  made  by  placing  a  few  pieces  of  phos- 
phorus in  the  bottom  of  a  good-sized  bottle  and  partly 
covering  them  with  water ;  and,  finally,  it  is  made  by 
treating  barium  dioxide,  BaO2,  and  some  other  compounds 
rich  in  oxygen  with  sulphuric  acid. 

Properties. — Ozone  is  a  gas  which  can  be  condensed  to 
the  liquid  form,  the  liquid  having  a  blue  color.  It  has  a 
strong  odor,  and  acts  in  an  irritating  way  upon  the  mem- 
branes lining  the  throat.  Its  chemical  conduct  is  entirely 
different  from  that  of  oxygen.  While  the  latter  at  ordi- 
nary temperatures  is  not  an  active  element,  ozone  is. 
It  acts  upon  most  substances  of  animal  or  vegetable 
origin,  and  oxidizes  nearly  all  the  metals,  besides  pro- 
ducing a  variety  of  other  changes  which  are  not  produced 
by  oxygen  at  the  ordinary  temperature.  Among  the 
characteristic  changes  which  may  be  made  use  of  for  the 
purpose  of  detecting  ozone  are  the  following :  (1)  It 
liberates  iodine  when  brought  in  contact  with  potassium 
iodide  ;  the  action  being  represented  thus  : 

2KI  +  H2O  +  O  =  2KOH  +  21. 

Now,  iodine  has  the  power  to  turn  starch  blue.  There- 
fore, if  ozone  is  brought  into  a  solution  containing  starch 
and  potassium  iodide,  a  blue  color  is  produced.  (2)  Ozone 
combines  with  metallic  silver  to  form  silver  peroxide, 
which  is  brown.  In  order  to  detect  ozone,  therefore,  a 
strip  of  polished  silver  may  be  exposed  to  the  gas,  and  if 
it  turns  brown  ozone  is  present. 

When  heated  to  300°  ozone  loses  its  characteristic  odor 
and  is  converted  into  ordinary  oxygen.  It  thus  appears 
that  we  can  start  with  ordinary  oxygen,  and  by  means  of 
an  electric  current  convert  it  into  a  substance  with  much 


RELATION  BETWEEN  OXYGEN  AND  OZONE. 


more  active  chemical  properties  and  differing  from  it  so 
markedly  that  one  would  hardly  suspect  the  close  relation 
between  the  two ;  and  then,  further,  by  simply  passing 
the  active  substance,  or  ozone,  through  a  tube  heated  to 
300°,  it  is  converted  back  again  into  oxygen  without  loss 
of  weight. 

Relation  between  Oxygen  and  Ozone. — When  experi- 
ment had  shown  that  oxygen  and  ozone  are  convertible 
one  into  the  other  without  change  of  weight,  the  suggestion 
was  made  that  the  difference  between  them  might  be  due 
to  a  difference  in  the  number  of  atoms  of  oxygen  contained 
in  the  molecule  of  each.  It  might  be,  for  example,  that 
in  the  molecule  of  oxygen  there  are  two  atoms  and  in  that 
of  ozone  three  or  more.  If  this  view  is  correct  there 
should  be  a  difference  between  the  specific  gravities  of 
the  two  gases,  and  by  a  study  of  this  difference  it  should 
be  possible  to  draw  a  conclusion  as  to  the  constitution  of 
the  molecule  of  ozone.  That  there  is  a  change  of  volume 
when  oxygen  is  changed  to  ozone  was  shown  by  enclosing 
the  former  in  a  tube  constructed  as  shown  in  Fig.  5.  The 
large  part  of  the  tube  is  furnished  with  two  platinum 
wires  which  pass  through  the  glass.  By  means  of  these 
wires  a  silent  discharge  of  electricity  is 
kept  up  through  the  oxygen,  and  it  is  thus 
partly  converted  into  ozone.  In  the  smaller 
bent  part  of  the  tube  there  is  a  small  col- 
umn of  concentrated  sulphuric  acid  which 
serves  as  a  stopper.  Any  change  in  the 
volume  of  the  gas  in  the  tube  will  cause 
this  sulphuric  acid  to  change  its  position. 
Now,  during  the  conversion  of  the  oxygen 
into  ozone  it  was  noticed  that  the  volume 
of  the  gas  decreased,  and  that,  on  heating 
the  tube  and  thus  converting  the  ozone 
into  oxygen  again,  the  original  volume  of 
the  latter  was  restored.  Unfortunately 
for  this  purpose  it  is  not  possible  by  any 
means  to  convert  more  than  a  comparatively 
small  proportion  of  the  oxygen  into  ozone, 
and  it  is  not  possible  in  the  experiment  described  to  de- 


FIG.  5. 


88  INORGANIC  CHEMISTRY. 

termine  the  relation  between  the  decrease  in  volume  and 
the  extent  of  the  conversion. 

Certain  ethereal  oils,  as  oil  of  turpentine,  have  the 
power  to  absorb  ozone  without  decomposing  it.  This 
fact  has  been  taken  advantage  of  for  the  purpose  of  de- 
termining the  constitution  of  ozone.  The  method  con- 
sisted in  ozonizing  oxygen,  heating  some  of  the  mixture 
and  measuring  the  increase  in  volume  caused  by  conver- 
sion into  oxygen  ;  and  then  treating  some  of  the  same 
mixture  with  oil  of  turpentine,  and  observing  the  de- 
crease in  volume.  By  this  means  it  is  possible  to  cal- 
culate the  change  in  volume  which  takes  place  when 
ozone  is  converted  into  oxygen.  The  results  showed 
that  two  volumes  of  ozone  give  three  volumes  of  oxygen, 
or  when  a  given  volume  of  oxygen  is  converted  into 
ozone  there  is  contraction  to  two-thirds  the  original  vol- 
ume. According  to  this,  the  molecular  weight  of  ozone 
by  Avogadro's  law  must  be  about  48,  for  from  3  volumes 
of  oxygen  there  are  formed  2  volumes  of  ozone,  which, 
interpreted  according  to  the  molecular  theory,  means 
that  three  molecules  of  oxygen  must  be  converted  into 
two  of  ozone ;  but  3  molecules  of  oxygen  weigh  96  parts  ; 
therefore,  2  molecules  weighing  the  same,  the  molecular 
weight  of  ozone  is  48.  The  further  conclusion  is  justified 
that,  while  the  molecule  of  ordinary  oxygen  consists  oi 
two  atoms,  that  of  ozone  consists  of  three  atoms.  The 
difference  may  be  represented  by  the  formulas  O2  or 

O 

A 
O=O  for  oxygen,  and  O3  or  O-O  for  ozone. 

Ozone  in  the  Air. — No  satisfactory  evidence  of  the  oc- 
currence of  ozone  in  the  air  has  ever  been  furnished.  Yery 
little  if  anything  is  known  in  regard  to  the  effect  of  ozone 
on  the  health.  Larger  quantities  would  undoubtedly  act 
injuriously.  It  is  commonly  believed  that  small  quanti- 
ties are  advantageous,  as  it  tends  to  destroy  substances 
which  are  unwholesome.  This  subject  has  not,  however, 
been  studied  with  sufficient  care  to  justify  a  positive 
opinion  in  regard  to  it.  The  difficulty  in  drawing  a  con- 
clusion is  increased  by  the  fact  that  there  are  other  sub- 


ALLOTROPT.  89 

stances  present  in  the  air  which  resemble  ozone  in  some 
respects,  as,  for  example,  hydrogen  dioxide,  H2Oa  (which 
see). 

Allotropy. — The  occurrence  of  an  element  in  two  or 
more  different  modifications  is  called  cdlotropy.  Thus, 
ozone  is  called  an  allotropic  form  of  oxygen.  There  are, 
however,  a  number  of  other  cases  of  a  similar  kind. 
Phosphorus,  for  example,  presents  itself  in  three  or  four 
different  varieties  which,  in  some  respects,  differ  mark- 
edly from  one  another.  Carbon  also  appears  in  three 
different  forms.  Whether  in  all  cases  the  difference  be- 
tween the  allotropic  forms  of  an  element  is  due  to  a  dif- 
ference in  the  molecular  constitution,  as  in  the  case  of 
oxygen,  is  impossible  to  decide  at  present,  for  the 
reason  that  the  molecular  weights  of  the  substances  can- 
not always  be  determined.  The  facts  learned  in  regard  to 
the  relations  between  oxygen  and  ozone  make  it  appear 
quite  probable  that  the  explanation  which  holds  good  for 
this  case  will  probably  hold  good  for  others. 

Varying  Number  of  Atoms  in  the  Molecules  of  one  and 
the  same  Element. — It  has  been  pointed  out  that  the  mol- 
ecules of  hydrogen,  oxygen,  and  some  other  elements 
consist  of  two  atoms  each.  We  have  just  seen  that  the 
molecule  of  ozone  consists  of  three  atoms.  There  are 
some  elements  wrhich  contain  a  different  number  of  atoms 
in  their  molecules,  according  to  the  temperature.  Thus, 
sulphur  is  a  solid  substance  which  boils  at  440°  C.  The 
specific  gravity  of  its  vapor  between  450°  and  500°  leads 
to  the  molecular  weight  about  192.  But  the  atomic 
weight  of  sulphur  is  very  nearly  32,  therefore  the  mole- 
cule of  sulphur  just  above  its  boiling-point  would  appear 
to  be  made  up  of  six  atoms.  As  the  temperature  is 
raised,  the  specific  gravity  becomes  less,  and  above  800° 
it  corresponds  to  the  molecular  weight  64,  showing  that 
at  this  high  temperature  the  molecule  apparently  con- 
sists of  two  atoms.  Above  that  point  the  specific  gravity 
does  not  change  materially.  According  to  the  latest 
investigations  on  this  subject,  it  appears  probable  that 
the  higher  specific  gravity  of  sulphur  vapor  at  the 


90  INORGANIC  CHEMISTRY. 

lower  temperature  is  to  be  ascribed  to  imperfect  vaporiza- 
tion of  the  substance,  and  that  the  evidence  that  the 
molecule  is  more  complicated  at  lower  than  at  higher 
temperature  is  not  conclusive.  Facts  somewhat  similar  to 
those  just  mentioned  have  been  brought  to  light  in  the  case 
of  iodine,  only  here  the  vapor  at  the  lower  temperature  has 
a  specific  gravity  showing  that  the  molecule  consists  of  two 
atoms,  but  at  very  high  temperatures  the  specific  gravity 
shows  that  the  molecule  contains  only  one  atom.  The 
atomic  weight  of  mercury  is  the  same  as  its  molecular 
weight,  or,  in  other  words,  the  molecule  and  atom  are  in 
this  case  identical. 

Nascent  State. — One  of  the  most  curious  phenomena 
connected  with  chemical  action  is  that  of  the  nascent 
state.  Some  elements,  which  under  ordinary  circum- 
stances are  inactive,  show  themselves  possessed  of  marked 
activity  if  they  are  allowed  to  act  the  instant  they  are  set 
free  from  their  compounds.  Thus,  if  carbon  monoxide, 
CO,  and  ordinary  oxygen  are  brought  together  at  the  or- 
dinary temperature,  no  action  takes  place.  If,  however, 
carbon  monoxide  is  brought  in  contact  with  something 
which  easily  gives  off  oxygen,  it  is  converted  into  carbon 
dioxide.  At  ordinary  temperatures,  for  example,  chro- 
mium trioxide  gives  up  oxygen  to  carbon  monoxide  as 
represented  thus  : 

SCO  +  2CrO3  =  3CO2  +  Cr2O3. 

So,  too,  hydrogen  at  ordinary  temperatures  is  an  inactive 
element,  but,  if  brought  in  contact  with  substances  the 
instant  it  is  set  free  from  a  compound,  it  produces  many 
marked  changes.  Many  substances  which  would  be  left 
unchanged,  if  hydrogen  were  passed  over  them  at  ordinary 
temperatures,  are  changed  if  put  in  the  liquid  from  which 
the  hydrogen  is  being  evolved.  The  most  plausible  ex- 
planation of  facts  like  these  is  to  be  found  in  the  molec- 
ular theory.  Free  oxygen  and  free  hydrogen  consist  of 
molecules  made  up  of  atoms  which  are  in  combination 
with  each  other.  Before  these  molecules  can  act  chemi- 
cally they  must  be  broken  up  into  atoms.  When  carbon 


HYDROGEN  DIOXIDE  OR  HYDROGEN  PEROXIDE.       91 

monoxide  is  brought  together  with  oxygen  in  the  mole- 
cular state,  the  condition  is  represented  thus  : 

co  +  o, 

Before  combination  can  take  place  the  molecule,  O2, 
must  be  decomposed  as  represented  thus  : 

0,  =  0  +  O  ; 

and  the  atoms  of  oxygen  then  combine  with  the  carbon 
monoxide.  Now  there  are  some  substances  which  yield 
oxygen  atoms  more  readily  than  ordinary  oxygen  does. 
This  is  true  of  chromium  trioxide,  CrO3,  which  in  con- 
tact with  substances  that  have  the  power  to  combine 
with  oxygen  breaks  up  thus : 

2CrO3  =  Cr,O3  +  O  +  O  +  O. 

When  hydrogen  is  liberated  from  a  compound  by 
chemical  action  the  uncombined  atoms  are  believed  to  be 
given  off  first ;  if  there  is  nothing  present  with  which 
the  hydrogen  can  combine,  these  atoms  combine  with 
each  other  in  pairs  to  form  the  molecules  of  ordinary 
hydrogen. 

The  examples  furnished  by  allotropism  and  the  phe- 
nomena of  the  nascent  state  show  how  chemical  facts 
which  otherwise  appear  entirely  incomprehensible  are 
satisfactorily  explained  by  the  aid  of  the  molecular 
theory. 

Hydrogen  Dioxide  or  Hydrogen  Peroxide,  H2O2. — When 
barium  dioxide  is  treated  with  hydrochloric  acid  or  di- 
lute sulphuric  acid  the  reactions  represented  by  the  fol- 
lowing equations  take  place : 

BaO2  +  2HC1    =  BaCl2   +  H2O2 ;  and 
Ba02  +  H2S04  =  BaS04  +  H2O2. 

In  the  latter  case  the  compound  BaSO4,  or  barium 
sulphate,  is  insoluble,  and  the  liquid  can  easily  be  sep- 
arated from  it  by  filtration.  If  the  liquid  is  boiled, 
decomposition  of  the  hydrogen  dioxide  takes  place,  oxy- 


92  INORGANIC  CHEMISTRY. 

gen  being  liberated  and  water  left  behind.     The  decom- 
position is  represented  thus  : 

2Ha02  =  2H20  +  Oa ;     or    H3O9  =  H,O  +  O. 

The  dioxide  can  be  obtained  in  the  form  of  a  color- 
less liquid  by  allowing  the  solution  to  stand  in  a  vacuum 
over  sulphuric  acid,  or  by  distilling  its  solution  in 
a  vacuum.  It  boils  at  84°— 85°  under  a  pressure  of 
68  mm. 

Properties. — Hydrogen  dioxide  is  a  clear,  syrupy  liquid 
which  does  not  solidify  at  —  30°.  It  is  characterized  by 
marked  instability.  It  breaks  down  slowly  even  at  or- 
dinary temperatures  if  simply  allowed  to  stand.  The  de- 
composition takes  place  easily  under  the  influence  of 
heat,  and  if  heated  rapidly  to  100°  explosion  is  apt  to 
take  place.  The  products  are  water  and  oxygen,  as  indi- 
cated above.  In  consequence  of  the  ease  with  which  hy- 
drogen dioxide  gives  off  oxygen  it  is  a  good  oxidizing 
agent,  and  it  is  now  manufactured  on  the  large  scale  for 
use  in  bleaching  and  in  medicine,  and  comes  into  the 
market  in  solutions  of  different  strengths.  The  solution 
has  a  bitter  taste  ;  if  concentrated,  it  affects  the  skin, 
causing  a  pricking  sensation,  and  making  white  spots, 
which,  however,  disappear  in  a  few  hours.  As  hydrogen 
dioxide  cannot  be  converted  into  vapor,  its  molecular 
weight  cannot  be  determined  by  the  method  of  Avo- 
gadro. 

Determinations  made  by  other  methods,  depending 
upon  observations  on  the  freezing-point  of  its  solutions 
(see  Chapter  XXIII),  show  that  its  molecule  is  repre- 
sented by  the  formula  H2O2 ,  and  not  by  HO.  Assuming 
that  the  formula  is  H2O2 ,  what  is  the  constitution  ?  "We 
have  very  little  clear  evidence  on  this  point,  but  the  most 
prevalent  view  is  that  the  atoms  are  connected  as  repre- 
sented in  the  formula  H-O-O-H,  in  which  the  oxygen 
atoms  are  in  combination  with  each  other,  and  also  with 
hydrogen.  The  principal  ground  upon  which  this  con- 
ception rests  is  that  this  is  the  only  way  in  which  two 
hydrogen  atoms  and  two  oxygen  atoms  can  be  joined  to- 


HYDROGEN  DIOXIDE— CHARACTERISTIC  REACTIONS.    93 

gether,  if  the  oxygen  atoms  are  bivalent  and  the  hydro- 
gen atoms  univalent.  This,  clearly,  is  not  good  evidence, 
as  it  is  quite  possible  that  oxygen  may  not  be  bivalent  in 
this  compound.  Indeed,  there  are  some  facts  which 
seem  to  show  pretty  conclusively  that  the  oxygen  atoms 
are  quadrivalent,  and  that  the  constitution  of  hydrogen 
dioxide  should  be  represented  by  the  formula 

H-0  =  0-H. 

It  is  impossible  here  to  discuss  the  significance  of  this 
formula.  Suffice  it  to  say  that  it  is  largely  based  upon 
the  results  of  observations  on  the  refractive  power  of  the 
compound. 

Occurrence  in  the  Air. — That  hydrogen  dioxide  occurs 
in  the  air  has  already  been  stated.  It  is  also  found  in 
rain  and  snow.  The  quantity  in  the  air  is  extremely 
small,  and  it  varies  at  different  times  of  the  day,  the 
.action  of  sunlight  being  evidently  favorable  to  its  forma- 
tion. 

Characteristic  Reactions. — Like  ozone,  hydrogen  diox- 
ide decomposes  potassium  iodide,  setting  iodine  free  : 

H2O2  +  2KI  =  2KOH  +  I2. 

This  fact  may  be  utilized  for  the  purpose  of  detecting 
the  compound.  The  separation  of  the  iodine  does  not 
take  place  readily  as  in  the  case  of  ozone,  but  the  action 
is  hastened  by  the  addition  of  a  very  small  quantity  of 
.a  dilute  solution  of  ferrous  sulphate,  FeSO4. — An  acid 
solution  of  potassium  permanganate  is  decolorized  by 
Jiydrogen  dioxide. — If  in  a  glass  cylinder  a  layer  of  ether 
is  poured  upon  a  solution  of  hydrogen  dioxide  and 
a  drop  of  a  solution  of  potassium  dichromate  is  then 
added,  and  the  cylinder  thoroughly  shaken,  the  ether 
will  take  up  a  blue  compound,  and  will  itself  become 
blue. 

When  hydrogen  dioxide  is  brought  together  with  sub- 
stances which  give  up  oxygen  readily,  action  generally 
takes  place  involving  decomposition  of  the  hydrogen 
•dioxide  as  well  as  of  the  other  substance.  Thus,  when 


94  INORGANIC  CHEMISTRY. 

it  is  brought  together  with  silver  oxide,  AgaO,  this  reac- 
tion takes  place  : 


So,  also,  it  undergoes  decomposition  with  ozone  as  rep- 
resented thus  : 

O,  +  HA  =  H,0  +  20,. 

The  explanation  of  these  facts  is  to  be  found  partly  in 
the  attraction  of  the  atoms  of  oxygen  for  each  other. 
In  the  molecule  of  silver  oxide  and  of  hydrogen  dioxide 
there  is  an  atom  of  oxygen  which  is  held  loosely.  When 
the  substances  are  brought  together  these  loosely  com- 
bined atoms  attract  and  combine  with  each  other,  as 
may  be  represented  thus  : 

(0,)0  +  0(H,0)  =  O,  +  O,  +  H,0. 

Thermochemical  Considerations.  —  By  methods  which 
need  not  be  described  here,  it  has  been  determined  that 
when  ozone  is  converted  into  oxygen  heat  is  evolved, 
and  the  thermochemical  equation  expressing  the  facts  is 
this  : 

2O3  =  3O2  =  +  2  X  36,200  cal. 

In  accordance  with  the  explanation  given  on  page  37, 
this  means  that  when  two  molecules  of  ozone  are  con- 
verted into  three  molecules  of  oxygen  72,400  c.  are 
evolved.  So,  also,  when  oxygen  .is  converted  into  ozone 
heat  is  absorbed,  the  equation  being 

(O2,  O)  =  O3  =  -  36,200  cal. 

To  convert  oxygen  into  ozone  therefore  requires  an 
addition  of  energy.  A  reaction  which  requires  an  addi- 
tion of  heat  is  called  an  endothermic  reaction,  and  one 
which  takes  place  with  an  evolution  of  heat  is  called  an 
exothermic  reaction.  In  general,  that  exothermic  reaction 
which  is  accompanied  by  the  greatest  evolution  of  heat 
takes  place  most  readily,  and  endothermic  reactions  do 
not  take  place  without  the  addition  of  energy  from  with- 


THERMOCHEMICAL   CONSIDERATIONS.  95 

out.  In  the  language  of  physics,  we  say  that  ozone  con- 
tains more  energy  than  oxygen,  and  therefore  it  acts 
more  readily. 

Somewhat  similar  relations  are  observed  between  hy- 
drogen dioxide  and  water.  The  decomposition  of  the 
dioxide  into  water  and  oxygen  is  accompanied  by  an  evo- 
lution of  heat : 

HaO3  =  HaO  +  0 ;  =  +  23,100  cal. ; 

and  the  formation  of  the  dioxide  requires  the  addition  of 
the  same  amount  of  heat : 

(H2O,  O)  =  -  23,100  cal. 

In  order  to  get  the  dioxide,  therefore,  the  reaction  must 
be  of  such  a  character  as  to  furnish  this  amount  of  heat. 
In  the  action  of  hydrochloric  acid  upon  barium  dioxide 
there  is  more  heat  evolved  than  is  required  in  the  forma- 
tion of  hydrogen  dioxide,  and  therefore  the  formation  in 
this  way  is  possible. 


CHAPTER  VIII. 

CHLORINE— HYDROCHLORIC  ACID. 

Historical. — Sodium  chloride  or  common  salt,  which  is 
the  principal  chlorine  compound  found  in  nature,  has- 
been  known  for  a  very  long  time.  In  1774  Scheele  first 
called  attention  to  chlorine  in  his  treatise  on  the  black 
oxide  of  manganese  or  manganese  dioxide.  In  accord- 
ance with  the  ideas  then  prevailing,  he  called  it  dephlo- 
gisticated  muriatic  acid.  Berthollet  suggested  in  1785 
that  it  was  oxidized  hydrochloric  acid,  and  it  was  then 
regarded  as  consisting  of  the  hypothetical  element,  mu- 
rium,  in  combination  with  oxygen.  In  1810  Davy  pointed 
out  that  the  idea,  previously  expressed  by  Gay-Lussac 
and  Thenard,  that  the  substance  is  an  element,  is  in  the 
highest  degree  probable,  and  he  gave  it  the  name  chlorine 
(from  XXoopos,  greenish-yellow).  Since  that  time  every- 
thing learned  in  regard  to  chlorine  has  gone  to  show  that 
it  is  an  element. 

Occurrence  of  Chlorine. — Though  widely  distributed  in 
nature,  chlorine  never  occurs  in  the  uncombined  state,  for 
the  reason  that  it  combines  with  other  substances  with 
great  ease,  and,  if  it  were  set  free,  it  would  at  once  enter 
into  combination.  It  does  not  occur  in  very  large  quantity 
as  compared  with  oxygen  and  hydrogen.  It  is  found 
chiefly  in  combination  with  the  element  sodium  as  common 
salt,  or  sodium  chloride,  a  compound  of  the  composition 
represented  by  the  formula  NaCl.  It  is  also  found  in 
combination  with  other  elements,  as  potassium,  magne- 
sium, etc.,  as  in  the  celebrated  mines  at  Stassfurt,  Ger- 
many. In  comparatively  small  quantity  it  occurs  in 
combination  with  silver,  forming  one  of  the  most  valua- 
ble silver  ores.  All  the  chlorine  which  we  have  to  deal 
with  is  made  from  common  salt. 

(96) 


PREPARATION  OF  CHLORINE.  97 

Preparation. — The  problem  to  be  solved  in  the  prepa- 
ration of  chlorine  from  common  salt  is  the  separation 
of  the  two  elements  sodium  and  chlorine.  This  cannot 
be  accomplished  directly  as  the  separation  of  mercury 
and  oxygen  in  the  decomposition  of  mercuric  oxide,  HgO, 
and  there  is  no  easily  obtained  compound  which  gives  off 
chlorine  when  heated.  The  method  adopted  consists  in 
making  hydrochloric  acid,  HC1,  from  sodium  chloride,  and 
then  treating  this  acid  with  some  substance  which  readily 
gives  off  oxygen.  The  change  of  sodium  chloride  to  hydro- 
chloric acid  is  readily  accomplished  by  treating  salt  with 
ordinary  sulphuric  acid — a  reaction  carried  on  on  the 
large  scale  in  the  manufacture  of  sodium  carbonate  or 
"  soda."  When  the  two  are  brought  together  a  change 
takes  place  which  will  be  studied  more  in  detail  farther 
on.  The  reaction  is  represented  by  the  equation 

(1)     2NaCl  +  H2SO4  =  Na2SO4  +  2HC1. 

Sodium  Sulphuric          Sodium         Hydrochloric 

chloride  acid  sulphate  acid 

As  will  be  seen,  the  sodium  of  the  sodium  chloride  and 
the  hydrogen  of  the  sulphuric  acid  exchange  places — a 
kind  of  action  which  is  quite  common. 

The  decomposition  of  the  hydrochloric  acid  and  libera- 
tion of  chlorine  under  the  influence  of  oxygen  takes  place 
as  represented  in  this  equation  : 

(2)    2HCl  +  O  =  HaO  +  Cl2. 

As  there  is  an  unlimited  supply  of  oxygen  in  the  air,  it 
would  be  advantageous  if  the  decomposition  of  the  hydro- 
chloric acid  could  be  effected  by  means  of  the  element  in 
the  free  state.  But  free  oxygen  alone  will  not  accom- 
plish the  change.  A  process  has  been  invented,  how- 
ever, for  the  manufacture  of  chlorine  on  the  large  scale 
which  depends  upon  the  decomposition  of  hydrochloric 
acid  by  the  oxygen  of  the  air.  This  is  Deacon's  process. 
It  consists  in  passing  hydrochloric  acid  and  air  together 
through  a  heated  tube  containing  clay  balls  saturated 
with  a  solution  of  copper  sulphate,  and  then  dried.  If 


98  INORGANIC  CHEMISTRY. 

the  temperature  of  the  tube  is  not  raised  too  high  the 
copper  sulphate  remains  unchanged.  Exactly  why  the 
oxidation  of  the  hydrochloric  acid  takes  place  under 
these  circumstances  is  not  positively  known,  but  it  prob- 
ably depends  upon  the  formation  and  decomposition  of 
intermediate  products.  Deacon's  process  is  used  quits 
extensively  on  the  continent  of  Europe,  while  in  England 
and  Scotland  another  method,  apparently  more  compli- 
cated, known  as  Weldon's  process,  is  chiefly  used.  In  the 
laboratory  chlorine  is  generally  made  by  treating  hydro- 
chloric acid  with  manganese  dioxide.  The  reaction  is 
represented  thus : 

(1)    MnO2  +  4HCl  =  MnCl2  +  2H2O  +  Cl2. 

This  is  explained  by  the  tendency  of  hydrogen  to  com- 
bine with  oxygen  to  form  water.  When  the  compound 
MnO  is  treated  with  hydrochloric  acid,  this  reaction 
takes  place  : 

MnO  +  2HC1  =  MnCl2  +  H2O. 

In  this  case  there  is  a  simple  exchange  of  places  by  the 
manganese  and  hydrogen  and  the  oxygen  and  chlorine, 
the  great  affinity  of  hydrogen  for  oxygen  being  a  promi- 
nent cause  of  the  change. 

So,  also,  when  manganese  dioxide  is  treated  with  hy- 
drochloric acid,  the  oxygen  may  be  first  replaced  by 
chlorine,  as  represented  in  the  equation 

MnO2  +  4HC1  =  MnCl4  +  2H2O. 

But  the  compound  MnCl4  gives  up  half  its  chlorine  when 
heated : 

MnCl4  =  MnCl2  +  C12 ; 

so  that  the  action  of  hydrochloric  acid  on  manganese 
dioxide  is  represented  by  the  equation  (1)  above.  Some 
recent  investigations  make  it  appear  probable  that  the 
reaction  is  more  complicated  than  it  is  here  repre- 
sented to  be ;  that  the  first  product  of  the  action  of 
hydrochloric  acid  on  manganese  dioxide  is  a  compound 


CHLORINE— WELDON'S  PROCESS.  99 

of  the  composition  H2MnCl6  (or  H2MnClB) ;  and  that  this 
breaks  down  tinder  the  influence  of  heat  thus : 

H2MnCl6  =  MnCl,  +  2HC1  +  C12. 

These  reactions  will  be  taken  up  under  the  head  of  Man- 
ganese (which  see). 

Instead  of  making  hydrochloric  acid  from  salt,  and 
then  treating  it  with  manganese  dioxide,  it  is  better  to 
mix  the  manganese  dioxide  and  common  salt  together, 
and  pour  upon  the  mixture  the  necessary  quantity  of 
sulphuric  acid.  In  this  case  the  manganese  dioxide  and 
sulphuric  acid  give  off  oxygen,  and  the  common  salt  and 
sulphuric  acid  give  off  hydrochloric  acid.  The  oxygen 
then  oxidizes  the  hydrochloric  acid,  and  chlorine  is  liber- 
ated. At  least  this  is  a  probable  explanation  of  the  reac- 
tion, for  it  is  known  that  when  manganese  dioxide  is 
heated  with  sulphuric  acid  oxygen  is  liberated,  as  repre- 
sented in  the  equation 

Mn02  +  H3S04  =  MnS04  +  H2O  +  O. 

Weldon's  Process. — As  there  is  a  large  demand  for 
chlorine,  much  attention  has  been  given  to  the  improve- 
ment of  the  methods  for  its  preparation.  One  of  the 
objections  to  the  ordinary  method  is  the  comparatively 
high  price  of  the  mineral,  manganese  dioxide.  As  this  is 
converted  into  the  chloride,  MnCl2,  in  the  preparation  of 
chlorine,  and  the  chloride  is  of  no  value,  the  expense  of 
preparation  is  quite  high.  A  process  has  been  invented 
for  the  regeneration  of  the  manganous  chloride,  MnCl,, 
or  for  the  conversion  of  this  compound  into  an  oxygen 
compound  which  with  hydrochloric  acid  will  give  chlo- 
rine. This  is  Weldon's  process.  It  will  be  taken  up 
under  the  head  of  Manganese  (which  see). 

Electrolytic  Process. — Various  processes  have  been 
devised  for  the  preparation  of  chlorine  by  the  action  of 
an  electric  current  on  a  solution  of  sodium  chloride  or  of 
hydrochloric  acid.  Such  electrolytic  processes  are  now 
in  use,  and  by  means  of  them  the  price  of  chlorine  has 
been  much  lowered.  1 0 

Properties.— Chlorine  is  a  greenish-yellow  gas.    It  has 


100  INORGANIC  CHEMISTRY. 

a  disagreeable  odor,  and  acts  upon  the  membranes  lining 
the  throat  and  nose,  causing  irritation  and  inflammation, 
suggesting  a  "  cold  in  the  head."  Inhaled  in  concen- 
trated form  it  causes  death.  Its  specific  gravity  at  20° 
is  2.48  (air  =  1),  and  as  compared  with  hydrogen  it  is 
35.18.  A  liter  of  chlorine  gas,  under  standard  conditions, 
weighs  3.162  grams.  It  is  soluble  in  water  and  acts 
upon  mercury,  and  therefore  cannot  be  collected  by  dis- 
placement of  either  of  these  liquids.  The  most  conven- 
ient way  to  collect  it  is  by  displacement  of  air.  It  can 
also  be  collected  over  warm  water  in  which  it  is  less 
soluble  than  in  cold  water,  or  over  a  saturated  solution 
of  sodium  chloride  in  which  it  is  but  slightly  soluble. 
It  is  easily  compressed  to  a  liquid,  and  is  now  sold  in 
this  form  in  cast-iron  cylinders. 

It  is  a  remarkably  active  substance,  combining  with 
or  acting  in  some  way  upon  most  other  substances  even 
at  ordinary  temperature.  This  activity  may  be  illus- 
trated by  introducing  into  vessels  containing  chlorine  a 
little  finely  powdered  antimony,  a  few  pieces  of  thin 
copper-foil,  a  piece  of  paper  with  ink-marks  on  it,  some 
flowers,  and  pieces  of  cotton-prints.  The  antimony  will 
take  fire  and  a  white  substance  will  be  formed.  The 
reaction  is  represented  by  the  following  equation  : 

Sb  +  301  =  Sb013. 

The  copper-foil  also  takes  fire,  and  is  converted  into  a 
chloride  as  represented  thus  : 

Cu  +  201  =  CuCla. 

Many  other  substances  unite  directly  with  chlorine 
with  evolution  of  heat  and  light,  and  form  compounds 
which  are  called  chlorides.  This  kind  of  action  is  of 
the  same  character  as  that  which  takes  place  in 
oxygen.  There  is,  however,  this  difference  between  re- 
actions in  oxygen  and  in  chlorine  :  the  latter  frequently 
take  place  at  ordinary  temperatures,  whereas  those  in 
oxygen  require  an  elevation  of  temperature  to  start 
them.  In  both  cases  the  gases  combine  with  the  other 


PROPERTIES  OF  CHLORINK^   /''•«>-»"  10£ 

substances  directly,  and  disappear  as  such,  and  the 
light  and  heat  are  caused  by  the  act  of  combination. 

Dry  liquid  chlorine,  at  its  boiling  temperature  (—33°.  6) 
does  not  act  on  potassium,  sodium,  or  aluminium. 

The  action  of  chlorine  upon  ink,  flowers,  and  cotton- 
print  illustrates  its  power  to  bleach.  It  is  important  to 
notice  that  if  the  colored  objects  be  introduced  dry  into 
dry  chlorine  the  action  does  not  take  place.  Moisture 
is  generally  essential  to  the  bleaching  by  chlorine. 
Chlorine  acts  directly  upon  some  dye-stuffs,  converting 
them  into  colorless  substances.  In  other  cases  it  has 
been  shown  that  the  destruction  of  the  color  is  due  to 
the  action  of  oxygen,  which  is  set  free  from  water  by 
chlorine.  In  direct  sunlight  chlorine  decomposes  water 
according  to  the  equation 


This  decomposition  can  be  illustrated  by  filling  a 
long  tube  with  a  solution  of  chlorine  in  water,  and  in- 
verting it  in  a  shallow  vessel  containing  some  of  the 
same  solution.  If  this  is  placed  in  the  direct  sunlight, 
bubbles  of  gas  will  be  seen  to  rise  in  the  tube  and  these 
will  collect  at  the  top,  while  the  color  of  the  solution, 
which  was  at  first  greenish-yellow  like  that  of  chlorine, 
will  disappear.  The  gas  which  collects  in  the  upper 
part  of  the  tube  is  oxygen. 

The  disintegrating  action  of  chlorine  upon  substances 
of  animal  and  vegetable  origin  may  be  illustrated  by 
moistening  a  piece  of  filter-paper  with  some  oil  of  turpen- 
tine, and  introducing  it  into  a  vessel  of  chlorine.  A  flash 
of  light  is  seen,  and  a  dense  black  cloud  is  formed. 
The  black  substance  is  mainly  carbon.  Oil  of  turpen- 
tine is  a  compound  of  carbon  and  hydrogen.  Chlorine 
abstracts  the  hydrogen  from  the  carbon,  leaving  the 
latter  mainly  in  the  uncombined  state.  If  the  chlorine 
is  allowed  to  act  slowly  upon  the  oil  of  turpentine  and 
similar  organic  substances,  the  chlorine  is  substituted 
atom  for  atom  for  the  hydrogen,  and  a  series  of  so-called 
substitution-products  is  obtained. 


"102  INORGANIC  CHEMISTRY. 

Chlorine  dissolves  readily  in  water  and  forms  a  solu- 
tion known  as  chlorine  water.  It  has  the  odor  and  color 
of  the  gas,  and  it  is  frequently  used  in  the  laboratory 
instead  of  the  gas.  From  what  has  been  said  it  is  evi- 
dent that  it  must  be  kept  protected  from  the  sunlight, 
or  decomposition  will  take  place,  resulting  in  the  forma- 
tion of  hydrochloric  acid  and  oxygen. 

Different  Kinds  of  Action. — A  careful  study  of  the  dif- 
ferent kinds  of  action  exhibited  by  chlorine  shows  that 
they  may  be  classified  under  three  heads  : 

(1)  First  it  acts  by  direct  combination  with  elements 
as  in  the  experiments  with  antimony  and  copper,  and,  as 
will  be  shown,  with  hydrogen  and  many  other  elements. 
Just  as  the  compounds  of  oxygen  with  other  elements 
are  called  oxides,  so  the  compounds  of  chlorine  with 
other  elements  are  called  chlorides.     Thus  the  compound 
of  antimony  and  chlorine,  SbCl3,  is  called  antimony  chlo- 
ride ;  that  of  zinc  and   chlorine,  ZnCl2,   is  called   zinc 
chloride  ;  etc.     In  case  an  element  forms  more  than  one 
compound  with  chlorine,  the  names  used  to  distinguish 
between   them   are   similar   to   those  used   for  oxides. 
Mercury  forms  two  chlorides  which  have  the  composi- 
tion represented  by  the  formulas  HgCl  and  HgCl2.    The 
one  with  the  smaller  proportion  of  chlorine  is  called 
mercurous  chloride ,  and  the  one  with  the  larger  propor- 
tion of  chlorine  is  called  mercuric  chloride.      So,  too, 
there  are  two  chlorides  of  iron  which  correspond  to  the 
formulas  FeCl2  and  FeCl3.     The  former  is  called  ferrous 
chloride,  and  the  latter  ferric  chloride. 

(2)  The  second  kind  of  action  of  chlorine  is  that  which 
is  called  substitution.     This  was  referred  to  in  connec- 
tion with  the  action  of  chlorine  on  the  oil  of  turpentine. 
The  general  character  of  this  kind  of  action  may  be  ex- 
plained by  the  aid  of  the  following  example.     There  is 
an  important  compound  of  carbon  and  hydrogen  called 
benzene,  which  has  the  formula  C6H6.     When  chlorine  is 
passed  through  this  compound,  which  is  a  liquid,  a  gas 
is  given  off  which  can  easily  be  shown  to  be  hydro- 
chloric, acid,  HC1.     This  action  continues  until  there  is 
no  hydrogen  left  in  combination  with  the  carbon,  but  in 


DIFFERENT  KINDS  OF  ACTION  OF  CHLORINE.     103 

place  of  the  benzene  there  is  now  a  compound  of  the 
formula  C6C16.  This  has  been  shown  to  be  the  final 
product  of  a  series  of  reactions  represented  by  the  fol- 
lowing equations  : 

C6H6  +  Cl,  =  C.HtCl  +HC1; 
C,H6C1  +  Cl,  =  C,H,Cla  +  HCl; 
C.H.C1,  +  Cl,  =  O.H.01.  +  HC1 ; 
C.H3C1S  +  Cl,  =  C.H,C14  +  HC1 ; 
CeH,Cl.  +  Cl,  =  C6HC1,  +  HC1 ; 
C6HC16  +  Cl,  =  C.C1,  +  HC1. 

In  each  stage  one  atom  of  chlorine  is  substituted  for  an 
atom  of  hydrogen,  but  the  hydrogen  does  not  escape  as 
such.  It  combines  with  chlorine  and  passes  off  in  the 
form  of  hydrochloric  acid. 

(3)  The  third  kind  of  action  is  that  noticed  in  bleach- 
ing, which  depends  upon  the  decomposition  of  water 
and  the  escape  of  oxygen  as  already  explained.  This 
action  does  not  take  place  in  the  dark,  but  does  take 
place  readily  in  the  direct  sunlight.  We  have,  however, 
seen  that  when  oxygen  acts  upon  hydrochloric  acid 
under  proper  conditions  water  is  formed  and  chlorine 
set  free.  It  appears,  therefore,  that,  under  some  cir- 
cumstances, this  reaction  is  possible  : 

2HC1  +  O  =  H2O  +  Cl, ; 
and,  under  other  circumstances,  this  one  : 
H2O  +  Cl,  =  2HC1  +  O. 

These  facts  appear  to  be  contradictory*  What  part  the 
sunlight  plays  is  not  known,  though  it  is  well  known  that 
it  is  capable  of  producing  a  great  variety  of  chemical 
changes.  We  shall  soon  see  that  it  is  only  necessary  to 
allow  it  to  act  for  an  instant  upon  a  mixture  of  hydrogen 
and  chlorine  to  cause  them  to  combine  with  violence. 
Then,  too,  the  various  processes  known  under  the  genera] 
name  of  photography  depend  upon  chemical  changes 
brought  about  by  the  sunlight.  Leaving  out  of  consid- 
eration this  kind  of  action,  the  decomposition  of  hydro- 
chloric  acid  by  oxygen  and  that  of  water  by  chlorine  can 


104  INORGANIC  CHEMISTRY. 

be  explained  by  a  consideration  of  the  heat  relations, 
The  heat  evoIVed  in  the  formation  of  1  molecule  of  water 
in  the  gaseous  form  is  58,069  cal.,  while  that  absorbed  in 
the  decomposition  of  2  molecules  of  hydrochloric  acid  is 
44,002  cal.  Therefore,  the  reaction, 

2H<21  +  O  =  H2O  +  01,, 

is  accompanied  by  Vn  evolution  of  heat.  It  is  exo- 
thermic, and  can  take  place  without  the  addition  of 
energy  from  without.  If  water  is  formed  as  a  liquid, 
the  heat  evolved  for  1  molecule  is  68,357  cal.,  while  that 
evolved  by  the  formation  of  2  molecules  of  hydrochloric 
acid  in  solution  is  78,630  cal.  Therefore,  the  heat 
evolved  in  the  formation  of  hydrochloric  acid  in  solu- 
tion is  greater  than  that  required  to  decompose  water, 
and  this  reaction  takes  place.  This  does  not,  however, 
explain  what  part  the  sunlight  plays  in  the  process. 

Chlorine  Hydrate  and  Liquid  Chlorine.  —  When  chlo- 
rine gas  is  passed  into  water  cooled  down  almost  to  the 
freezing-point,  crystals  appear  in  the  vessel.  These  con- 
sist of  chlorine  and  water  as  represented  by  the  formula 
Cl  -f-  5H2O ;  or,  assuming  that  it  is  formed  by  the  com- 
bination of  the  molecules  of  chlorine  with  water,  the 
formula  should  be  written  Cla  +  10H2O.  It  gives  off 
chlorine  at  the  ordinary  temperature  and,  if  allowed  to 
stand,  undergoes  complete  decomposition  into  chlorine 
and  water.  If  gently  heated  the  chlorine  is  given  off 
rapidly.  This  fact  was  taken  advantage  of  by  Faraday 
for  the  purpose  of  subjecting  the  gas  to  high  pressure 
and  low  temperature.  For  this  purpose  he  placed  some 
of  the  hydrate  in  a  strong  glass  tube  of  the  form  repre- 
sented in  Fig.  6.  The  com- 
pound was  put  in  the  part  of 
the  tube  marked  db,  and  the 
other  end,  c,  then  sealed.  The 
arm  ab  was  warmed  by  dip- 
ping it  in  warm  water,  while 
the  other  arm  was  placed  in 
Fl0-6-  a  freezing  mixture.  Under 

these  circumstances  the  chlorine  is  given  off  from  the 
hydrate,  but  being  unable  to  escape  from  the  tube  the 


HYDROCHLORIC  ACID.  105 

pressure  is  increased  to  such  an  extent  that  at  the  low 
temperature  the  gas  assumes  the  liquid  form. 

Applications  of  Chlorine. — Chlorine  is  used  very  exten- 
sively in  the  arts,  particularly  for  the  purpose  of  bleach- 
ing. It  is  also  used  for  the  manufacture  of  a  large  num- 
ber of  compounds  which  contain  chlorine,  the  principal 
ones  being  bleaching  powder  or  calcium  hypochlorite, 
and  potassium  chlorate.  If  used  in  sufficient  quantity 
chlorine  is  an  excellent  disinfectant  and  deodorizer. 
By  far  the  largest  quantity  of  the  chlorine  manufactured 
is  converted  into  bleaching  powder  or  calcium  hypochlo- 
rite, as  this  can  be  conveniently  transported,  and  the 
chlorine  can  be  obtained  from  it  when  desired.  It  is 
only  necessary  to  expose  it  to  the  air  to  effect  a  partial 
decomposition  accompanied  by  a  liberation  of  chlorine  ; 
and  the  addition  of  hydrochloric  or  sulphuric  acid  causes 
it  to  give  it  up  completely,  as  will  be  shown  farther  on. 
This  bleaching  powder  is  now  used  almost  exclusively 
instead  of  chlorine  gas  for  bleaching. 

HYDEOCHLOKIC  ACID. 

Historical. — Hydrochloric  acid  was  first  prepared  in 
large  quantity  by  Glauber  in  the  seventeenth  century, 
and  his  description  is  not  unlike  those  which  one  fre- 
quently reads  nowadays  referring  to  some  patent  medi- 
cine. The  method  of  preparation  used  by  him  was  the 
same  as  that  used  at  present,  viz.,  the  action  of  sulphuric 
acid  upon  common  salt. 

Study  of  the  Action  of  Hydrogen  upon  Chlorine. — If  hy- 
drogen and  chlorine  are  brought  together  in  the  dark  no 
action  takes  place,  no  matter  how  long  they  are  allowed 
to  stand  together.  If,  however,  the  mixture  is  put  in 
diffused  sunlight,  gradual  combination  takes  place ; 
and  if  the  direct  light  of  the  sun  is  allowed  to  shine 
for  an  instant  on  the  mixture,  explosion  occurs,  and  this 
is  the  sign  of  the  combination  of  the  two  gases.  The 
same  sudden  combination  is  effected  by  applying  a 
flame  or  spark  to  the  mixture,  or  by  illuminating  it  in- 


106  INORQANIC  CHEMISTRY. 

stantaneously  with  the  light  from  burning  magnesium 
or  an  electric  light.  On  comparing  these  facts  with 
those  learned  in  studying  the  action  of  hydrogen  on 
oxygen  a  marked  difference  is  evident.  Hydrogen  and 
oxygen  do  not  combine  either  in  the  dark  or  the  direct 
sunlight,  but  only  when  a  spark  is  brought  in  contact 
with  the  mixture. 

Another  way  in  which  hydrogen  can  be  made  to 
combine  with  chlorine  is  by  introducing  a  jet  of  burning 
hydrogen  into  a  vessel  containing  chlorine.  The  hydro- 
gen will  continue  to  burn,  but  the  character  of  the  flame 
will  change  completely,  and  above  the  vessel  white 
fumes  will  be  observed.  This  burning  of  hydrogen  in 
chlorine  is  entirely  analogous  to  the  burning  of  hydro- 
gen in  oxygen.  It  is  simply  an  act  of  combination  of 
the  two  gases,  in  each  case,  accompanied  by  an  evolution 
of  light  and  heat.  And  just  as  oxygen  can  be  burned  in 
hydrogen  by  a  proper  arrangement  of  apparatus,  so 
chlorine  can  also  be  burned  in  hydrogen. 

To  determine  the  relation  between  the  volumes  of  hy- 
drogen and  chlorine  which  combine  with  each  other  and 
the  volume  of  the  product  formed  is  more  difficult  than 
in  the  case  of  hydrogen  and  oxygen,  mainly  for  the  rea- 
son that  chlorine  acts  upon  mercury  and  is  dissolved  by 
water.  It  is  necessary  to  proceed  indirectly. 

Instead  of  causing  hydrogen  and  chlorine  to  combine, 
hydrochloric  acid  is  decomposed  and  the  volumes  of  the 
hydrogen  and  chlorine  obtained  are  determined.  One 
method  of  effecting  this  consists  in  decomposing  hydro- 
chloric acid  by  an  electric  current  in  an  apparatus 
like  that  referred  to  in  connection  with  the  decom- 
position of  water.  As  chlorine  is,  however,  soluble 
in  water,  the  apparatus  is  filled  with  a  saturated  solution 
of  common  salt  to  which  a  strong  solution  of  hydro- 
chloric acid  in  water  is  added.  On  passing  a  fairly 
strong  current,  the  hydrochloric  acid  is  decomposed, 
hydrogen  being  given  off  at  one  pole  and  chlorine  at 
the  other.  For  a  given  volume  of  hydrogen  the  same 
volume  of  chlorine  is  liberated,  which  makes  it  ap- 
pear probable  that  hydrogen  and  chlorine  are  combined 


PREPARATION  OF  HYDROCHLORIC  ACID.  107 

in  hydrochloric  acid  in  the  proportion  of  volume   to 
volume. 

For  the  purpose  of  further  studying  the  volume  re- 
lations, the  following  experiment  is  of  value.  A  tube  is 
filled  with  hydrochloric  acid  gas.  A  small  piece  of  po- 
tassium is  then  introduced,  when  decomposition  takes 
place  as  represented  in  the  equation 


The  gas  left  in  the  vessel  is  hydrogen,  as  can  easily  be 
shown.  On  measuring  its  volume  it  is  found  to  be  just 
half  that  of  the  hydrochloric  acid  gas  decomposed. 
Taking  this  fact  into  consideration  with  the  fact  that 
whenever  hydrochloric  acid  is  decomposed  by  an  electric 
current  equal  volumes  of  hydrogen  and  chlorine  are  ob- 
tained, it  appears  that  in  the  formation  of  hydrochloric 
acid  gas  1  volume  of  hydrogen  combines  with  1  volume 
of  chlorine  to  form  2  volumes  of  hydrochloric  acid,  a 
fact  which  was  referred  to  in  the  chapter  on  the  Atomic 
Theory.  The  weight  of  the  hydrogen  is  found  to  bear 
to  the  weight  of  the  hydrochloric  acid  the  proportion 
1  :  36.18.  In  other  words,  in  36.18  parts  of  hydro- 
chloric acid  there  are  35.18  parts  of  chlorine  and  1  part 
of  hydrogen. 

Preparation.  —  For  the  preparation  of  hydrochloric  acid 
in  the  laboratory  as  well  as  on  the  large  scale,  ordinary 
sulphuric  acid  is  poured  upon  common  salt.  Two  reac- 
tions may  take  place  between  these  substances,  depend- 
ing largely  upon  the  amount  of  sulphuric  acid  used.  If 
the  two  substances  are  brought  together  in  the  propor- 
tion of  the  weights  of  their  molecules  or  their  molecular 
weights,  the  principal  reaction  is  the  one  represented  in 
the  following  equation  : 

NaCl  +  H2S04  =  NaHS04  +  HC1. 

In  this  case  the  atom  of  sodium  of  the  molecule  of 
sodium  chloride  is  substituted  for  one  atom  of  hydrogen 
in  the  molecule  of  sulphuric  acid,  while  the  hydrogen 
and  chlorine  unite  to  form  hydrochloric  acid.  If,  on  the 
other  hand,  the  substances  are  brought  together  in  the 


108  INORGANIC  CHEMISTRY. 

proportion  of  2  molecules  of  sodium  chloride  and  1 
molecule  of  sulphuric  acid  the  principal  reaction  is  the 
following : 

2NaCl  +  H2S04  =  Na3SO4  +  2HC1. 

Properties. — Hydrochloric  acid  is  a  colorless  trans- 
parent gas,  and  has  a  sharp  penetrating  taste  and  smell. 
Inhaled  it  produces  suffocation.  It  is  extremely  easily 
soluble  in  water,  1  volume  of  water  at  ordinary  temper- 
atures dissolving  450  times  its  own  volume  of  the  gas,  and 
at  0°,  500  times.  The  solution  of  the  gas  in  water  is  what 
is  generally  called  hydrochloric  acid.  So  great  is  the  at- 
traction of  the  gas  for  water  that  it  condenses  moisture 
from  the  air  ;  hence,  although  the  gas  itself  is  quite  color- 
less and  transparent,  when  it  comes  in  contact  with  the  air 
dense  white  clouds  are  formed,  which  are  not  formed  if  it 
is  kept  from  contact  with  the  air,  as  can  easily  be  shown 
by  filling  glass  vessels  with  the  gas.  Hydrochloric  acid 
does  not  burn  and  does  not  support  combustion.  This 
is  equivalent  to  saying  that  it  does  not  combine  with 
oxygen  under  ordinary  circumstances,  and  that  sub- 
stances which  combine  with  the  oxygen  of  the  air  do  not 
combine  with  hydrochloric  acid.  On  the  other  hand,  we 
have  seen  that  under  some  circumstances  oxygen  does 
act  upon  hydrochloric  acid  and  cause  an  evolution  of 
chlorine.  The  gas  is  comparatively  easily  condensed  to 
the  liquid  form. 

When  a  concentrated  solution  of  hydrochloric  acid  in 
water  is  heated,  gas  is  given  off,  but  if  a  dilute  solution 
is  heated  water  is  given  off.  In  either  case,  when  the 
composition  of  the  liquid  is  that  represented  by  the 
formula  HC1  -f-  8H2O,  it  boils  under  the  ordinary  pres- 
sure of  the  atmosphere  unchanged.  If  the  pressure  is 
lowered  the  composition  of  the  liquid  which  passes  over 
in  the  process  of  distillation  changes,  so  that  it  contains 
a  larger  percentage  of  hydrochloric  acid  the  lower  the 
pressure  becomes.  This  fact  seems  to  show  that  the 
liquid  of  the  composition  HC1  +  8H2O,  which  boils 
unchanged  at  the  temperature  110°  under  the  ordinary 
pressure  of  the  atmosphere  is  not  a  chemical  compound. 


HYDROCHLORIC  ACID.  109 

On  the  other  hand,  it  certainly  does  not  conduct  itself 
like  most  ordinary  solutions  of  gases. 

There  is  a  definite  compound  of  hydrochloric  acid 
with  water  called  hydrochloric  acid  hydrate,  which  has 
the  composition  HC1  -|-  2H2O.  This  is  formed  by  pass- 
ing hydrochloric  acid  gas  into  the  concentrated  aqueous 
solution  cooled  down  to  —  22°.  Under  these  circum- 
stances the  hydrate  separates  in  the  form  of  crystals. 

Commercial  hydrochloric  acid  is  a  yellowish  liquid,  the 
color  being  due  to  the  presence  of  impurities,  such  as 
iron  and  organic  substances.  The  solution  is  obtained 
in  the  factories  in  which  "  soda  "  or  sodium  carbonate 
is  made.  This  is  an  extremely  important  substance  in 
the  arts.  It  does  not  occur  in  nature,  but  is  manu- 
factured from  common  salt.  In  the  process  most  com- 
monly used  salt  is  first  converted  into  sodium  sulphate, 
Na2SO4,  by  treating  it  with  sulphuric  acid.  Hydro- 
chloric acid  is  necessarily  given  off.  When  the  factories 
were  first  established  in  England,  the  gas  was  allowed  to 
escape  as  a  waste  product,  but  the  effects  produced  by 
it  upon  the  vegetation  of  the  surrounding  country  were 
so  injurious  that  a  law  was  passed  prohibiting  the 
manufacturers  from  allowing  the  gas  to  escape.  It  is 
now  collected  by  causing  it  to  pass  through  towers  so 
constructed  as  to  expose  a  large  surface  of  bricks  or 
sandstone  plates  over  which  a  current  of  cold  water  is 
constantly  kept  flowing.  This  water  dissolves  the  hy- 
drochloric acid,  and  the  solution  collected  below  is 
commercial  hydrochloric  acid.  In  this  way  enormous 
quantities  of  the  acid  are  produced,  but  its  uses  are  nu- 
merous and  it  always  commands  a  price. 

Pure  hydrochloric  acid  is  a  solution  of  the  pure  gas  in 
pure  water.  It  is  colorless,  and  when  concentrated  it 
gives  off  fumes  when  exposed  to  the  air.  The  solution 
when  heated  gives  off  a  large  part  of  the  gas  contained  in 
it,  and  by  boiling  it  can  all  be  evaporated. 

Chemical  Action  of  Hydrochloric  Acid. — If  the  action  of 
hydrochloric  acid  towards  the  elements  should  be  studied 
systematically  it  would  be  found  that  many  of  them 
act  by  simply  taking  the  place  of  the  hydrogen,  as  has 


110  INORGANIC  CHEMISTRY. 

already  been  illustrated  in  the  preparation  of  hydrogen 
by  treating  hydrochloric  acid  with  zinc  when  this  reac- 
tion takes  place : 

Zn  +  2HC1  =  ZnCl2  +  H2. 
With  iron  the  reaction  is  : 

Fe  +  2HC1  =  FeCl2  +  H3 ; 
with  tin : 

Sn  +  2HC1  =  SnCla  +  H2 ; 
with  potassium : 

K   +  HC1  '  =  KOI     +  H,  or 
Ka  +  2HC1  =  2KC1   +  H2 ; 

with  sodium : 

Naa  +  2HC1  =  2NaCl  +  H2 ; 

with  calcium : 

Ca  +  2HC1  =  Ca012  +  H3 ;  etc.,  etc. 

On  the  other  hand,  there  are  many  elements  which  do 
not  act  in  this  way  towards  hydrochloric  acid.  Sul- 
phur, nitrogen,  phosphorus,  carbon,  and  boron  may  be 
taken  as  examples.  These  elements  do  not  act  upon  hy- 
drochloric acid  at  all.  We  might,  therefore,  divide  the 
elements  into  two  classes  :  (1)  those  which  act  upon  hy- 
drochloric acid  setting  hydrogen  free  and  forming  chlo- 
rides ;  and  (2)  those  which  do  not  act  upon  hydrochloric 
acid. 

Again,  when  hydrochloric  acid  acts  upon  the  oxides  of 
those  elements  which  have  the  power  to  liberate  hydro- 
gen from  it,  it  forms  the  same  chlorides  as  are  formed 
when  the  element  alone  acts,  but  instead  of  hydrogen 
being  liberated  water  is  formed.  Thus  when  the  acid 
acts  upon  zinc  oxide,  ZnO,  the  reaction  takes  place  thus  : 

ZnO  +  2HC1  =  ZnCl2  +  H2O ; 


CHEMICAL  ACTION  OF  HYDROCHLORIC  ACID.     .111 
with  lime  or  calcium  oxide,  CaO,  it  is  : 

CaO  +  2HC1  =  CaCla  +  HaO  ; 
with  potassium  oxide  : 

K2O  +  2HC1  =  2KC1  +  H2O  ;  etc.,  etc. 

In  these  reactions  there  are  two  forces  at  work  tending 
to  effect  the  change.  There  is,  first,  the  affinity  of  the 
element,  which  is  combined  with  oxygen,  for  chlorine, 
and>  second,  the  affinity  of  the  hydrogen  for  oxygen. 
"We  should,  therefore,  naturally  expect  hydrochloric 
acid  to  act  more  readily  upon  the  oxides  than  upon  ele- 
ments, and  it  has  been  found  that  this  is  the  case. 

If  the  so-called  hydroxides  of  those  elements  which 
act  upon  hydrochloric  acid  be  brought  in  contact  with 
the  acid,  action  takes  place  even  more  readily  than  with 
the  oxides,  and  the  products  are  the  same.  Thus  potas- 
sium hydroxide  and  hydrochloric  acid  give  potassium 
chloride  and  water : 

KOH  +  HC1  =  KC1  +  H2O  ; 

calcium  hydroxide  and  hydrochloric  acid  give  calcium 
chloride  and  water,  thus  : 

Ca(OH),  +  2HC1  =  CaCla  +  2H2O  ; 

aluminium  hydroxide  and  hydrochloric  acid  give  alumin- 
ium chloride  and  water,  thus  : 

Al(OH),  +  3HC1  =  A1C1,  +  3H2O ;  etc.,  etc. 

There  are  then  elements  which  act  upon  hydrochloric 
acid  liberating  hydrogen  and  forming  chlorides  ;  and 
the  oxides  and  hydroxides  of  these  elements  act  upon 
hydrochloric  acid  forming  chlorides  and  water.  The 
elements  which  act  in  this  way  are  commonly  called 
metals  or,  for  reasons  which  will  be  discussed  farther  on, 
base-forming  elements. 

If  those  elements  which  do  not  set  hydrogen  free  from 


112  INORGANIC  CHEMISTRY. 

hydrochloric  acid  are  treated  directly  with  chlorine,  they 
generally  combine  with  it  to  form  chlorides.  But  these 
chlorides  differ  markedly  from  the  chlorides  of  the  met- 
als, especially  in  their  conduct  towards  water.  Two  ex- 
amples will  suffice  for  the  present.  Phosphorus  forms 
a  chloride  of  the  formula  PC13,  known  as  phosphorus- 
trichloride.  In  contact  with  water  it  undergoes  decom- 
position according  to  this  equation  : 

PC13  +  3H30  =  PO3H3  +  3HC1 ; 

so,  too,  the  chloride  of  boron,  BC13,  undergoes  the  same 
kind  of  change : 

BC13  +  3H2O  =  BO3H3  +  3HC1. 

Similarly,  the  other  chlorides  of  the  elements  of  this 
class  tend  to  pass  into  the  oxides  or  hydroxides  when 
brought  in  contact  with  water.  Those  elements  which 
do  not  act  upon  hydrochloric  acid  setting  hydrogen  free 
and  forming  chlorides  are  generally  called  non-metals  or, 
for  reasons  which  will  appear  later,  acid-forming  elements. 
The  chlorides  of  the  acid-forming  elements  are  generally 
decomposed  by  water  and  the  corresponding  oxides  or 
hydroxides  are  formed.  In  general  terms,  the  oxide 
of  a  base-forming  element  or  of  a  metal  is  acted  upon  by 
hydrochoric  acid,  a  chloride  and  water  being  formed ; 
and  the  chloride  of  an  acid-forming  element  or  of  a  non- 
metal  is  acted  upon  by  water,  a  hydroxide,  or  oxide,  and 
hydrochloric  acid  being  formed.  We  shall  have  many 
illustrations  of  the  opposite  chemical  character  of  these 
two  classes  of  elements,  and  we  shall  see  that  many  of 
the  most  important  and  characteristic  chemical  reac- 
tions are  associated  with  these  differences. 


CHAPTER  IX. 

COMPOUNDS  OF  CHLORINE  WITH  OXYGEN  AND  WITH 
HYDROGEN  AND  OXYGEN. 

General. — As  has  been  seen,  chlorine  combines  very 
readily  with  hydrogen,  and  hydrogen  with  oxygen,  and 
the  products  are  stable  compounds.  On  the  other 
hand,  chlorine  cannot  be  made  to  combine  directly  with 
oxygen.  By  indirect  processes  they  can  be  combined, 
but  the  compounds  undergo  decomposition  easily,  yield- 
ing back  the  chlorine  and  oxygen  contained  in  them. 
Before  taking  up  the  compounds  of  chlorine  and  oxygen, 
however,  it  will  be  best  to  discuss,  as  far  as  may  be 
necessary,  the  compounds  of  chlorine,  hydrogen,  and 
oxygen  which  are  more  easily  made,  and  from  which  the 
oxides  are  made. 

Principal  Reactions  for  making  Compounds  of  Chlorine 
with  Hydrogen  and  Oxygen. — One  of  the  principal  reac- 
tions made  use  of  for  the  preparation  of  compounds  of 
chlorine,  oxygen,  and  hydrogen  consists  in  treating  po- 
tassium hydroxide  with  chlorine.  The  strong  affinity  of 
chlorine  for  potassium  shown  by  the  decomposition  of 
hydrochloric  acid  by  potassium  would  lead  us  to  expect 
that  when  chlorine  acts  upon  potassium  hydroxide,  po- 
tassium chloride,  KC1,  would  be  formed : 

KOH  +  01  =  KC1  +  O  +  H. 

But  it  also  has  a  strong  affinity  for  hydrogen,  so  that 
hydrochloric  acid  would  be  formed  as  well  as  potassium 
chloride  : 

KOH  +  2C1  =  KC1  +  HC1  +  O. 

The  oxygen  can,  however,  combine  with  potassium 
chloride  and  form  compounds,  KC1O,  KC1O2,  KC1O3,  and 

(113) 


114  INORGANIC  CHEMISTRY. 

KC1O4 ;  and  hydrochloric  acid,  if  formed,  would  com> 
bine  with  potassium  hydroxide,  thus  : 

KOH  +  HC1  =  KC1  +  H2O. 

When  potassium  hydroxide  is  treated  with  chlorine 
we  may  therefore  expect  to  obtain  potassium  chloride, 
KC1 ;  some  compound  containing  potassium,  chlorine, 
and  oxygen ;  and  water.  Experiment  has  shown  that 
the  action  takes  place  as  we  should  expect,  and  that  the 
compound  of  potassium,  chlorine,  and  oxygen  is  differ- 
ent according  to  the  conditions  of  the  experiment.  If 
the  solution  of  caustic  potash  is  warm  and  concentrated 
the  product  is  richer  in  oxygen  than  when  the  solution 
is  dilute  and  cold.  With  the  concentrated  solution  the 
reaction  takes  place  thus  : 

6KOH  +  3C12  =  5KC1  +  KC1O3  +  3H2O. 

While  five  atoms  of  potassium  appear  in  the  form  of  the 
chloride,  one  appears  in  the  form  of  an  oxygen  com- 
pound, KC1O3,  potassium  chlorate,  which  we  have  already 
had  to  deal  with  in  connection  with  the  preparation  of 
oxygen.  With  the  dilute  solution  of  caustic  potash  the 
reaction  takes  place  thus  : 

2KOH  +  Cla  =  KC1  +  KC1O  +  H2O. 

The  oxygen  product  in  this  case  is  potassium  hypocMo- 
rite,  KC1O. 

Potassium  chlorate,  KC1O3,  and  potassium  hypocMorite, 
KC1O,  bear  the  same  relation  to  two  compounds,  HC1O3 
and  HC1O,  that  potassium  chloride,  KC1,  and  sodium 
chloride,  NaCl,  bear  to  hydrochloric  acid.  But  we  have 
seen  that  hydrochloric  acid  can  easily  be  obtained  from 
sodium  chloride  by  treating  it  with  sulphuric  acid. 
Potassium  chloride  undergoes  the  same  change  when 
treated  with  sulphuric  acid.  Indeed,  we  shall  see  that 
nearly  all  compounds  containing  sodium  or  potassium 
give  up  these  metals  when  treated  with  sulphuric  acid, 
and  take  up  hydrogen  in  the  place  of  them. 


CHLORIC  ACID.  115 

Treating  potassium  chloride  with  sulphuric  acid  this 
reaction  takes  place  : 

2KC1  +  H2S04  =  K2S04  +  2HCL 

Similarly,  treating  potassium  chlorate  with  sulphuric 
acid,  this  reaction  takes  place  : 

2KC1O3  +  HaSO4  =  K8SO4  +  2HC1O3. 

The  compound  HC1O3  is  called  cUoric  acid.  Further, 
when  potassium  hypochlorite  is  decomposed  by  sul- 
phuric acid  under  proper  circumstances,  it  undergoes 
the  same  kind  of  decomposition  : 

2KC10  +  H2SO4  =  K2SO4  +  2HC1O. 

The  compound  HC1O  is  called  hypochlorous  acid. 

Chloric  Acid,  HC1O3.— The  preparation  of  chloric  acid 
from  potassium  chlorate  is  accomplished  by  treating 
a  water  solution  of  the  chlorate  with  fluosilicic  acid, 
H2SiF6,  with  which  potassium  forms  an  insoluble  com- 
pound. The  reaction  which  takes  place  is  represented 
thus: 

2KC1O3  +  HaSiF6  =  K2SiF6  +  2HC1O3. 

The  compound  K2SiF6  is  known  as  potassium  fluosili- 
cate.  It  will  be  observed  that  this  reaction  is  of  the 
same  general  character  .as  that  represented  above  as  tak- 
ing place  between  potassium  chlorate  and  sulphuric  acid. 
The  difference  between  the  two  reactions  is  that  potas- 
sium fluosilicate  is  insoluble  in  water,  while  potassium 
sulphate  is  soluble.  By  using  fluosilicic  acid,  therefore, 
a  solution  of  chloric  acid  is  obtained  free  from  other  sub- 
stances, provided  just  enough  of  the  fluosilicic  acid  is 
added  to  form  potassium  fluosilicate  with  all  the  potas- 
sium. This  solution  is  unstable,  and  if  heated  above  40° 
the  chloric  acid  undergoes  decomposition  according  to 
the  equation 

4HC1O,  =  01,  +  3O  +  2HC1O4  +         " 


116  INORGANIC  CHEMISTRY. 

Properties. — Chloric  acid  acts  upon  metals  in  the  same 
general  way  that  hydrochloric  acid  does.  It  gives  up  its 
hydrogen  and  takes  up  metal  in  its  place  forming  com- 
pounds like  potassium  chlorate,  KC1O3,  sodium  chlorate, 
NaClO3,  etc.  In  consequence  of  the  ease  with  which  it 
gives  up  oxygen,  it  is  used  extensively  for  the  prepara- 
tion of  oxygen,  and  for  the  purpose  of  adding  oxygen  to 
other  substances,  or  as  an  oxidizing  agent.  Potassium 
chlorate  and  other  compounds  of  similar  character  de- 
rived from  chloric  acid  are  used  in  the  manufacture  of 
fire- works.* 

Hypochlorous  Acid,  HC1O. — The  formation  of  potas- 
sium hypochlorite,  KC1O,  by  treating  caustic  potash  with 
chlorine  has  been  mentioned.  A  similar  reaction  is  em- 
ployed on  the  large  scale  in  the  manufacture  of  bleach- 
ing powder  or  "  chloride  of  lime."  This  consists  in 
treating  slaked  lime  or  calcium  hydroxide  with  chlo- 
rine. The  action  is  represented  thus  : 

2Ca  (OH)2  +  2C12  =  Ca(ClO)2  +  CaCl2  +  2H2O. 

The  compound  Ca(ClO)2  known  as  calcium  hypochlorite 
is  derived  from  hypochlorous  acid  by  replacing  two 
atoms  of  hydrogen  in  two  molecules  of  the  acid  by  one 
atom  of  the  bivalent  metal  calcium 

;  ggjg  gives  Ca(gg)  or  Ca(ClO),.  .,  J 

Just  as  a  mixture  of  potassium  chloride  and  potassium 
hypochlorite  is  formed  when  potassium  hydroxide  is 
used,  so  apparently  a  mixture  of  calcium  chloride  and 
calcium  hypochlorite  is  formed  when  calcium  hydroxide 
is  used.  This  point  will  be  discussed  to  some  extent 
under  the  head  of  Calcium  Hypochlorite  (which  see), 
when  it  will  be  shown  that  there  are  good  reasons  for 


*  Great  care  is  necessary  when  working  with  potassium  chlorate,  as 
with  many  substances  it  forms  explosive  mixtures.  Treated  with  con- 
centrated  acids  it  undergoes  rapid  and  violent  decomposition. 


HYPOCHLOBOUS  ACID.  11? 

believing  that  the  product  called  bleaching  powder  is  a 
distinct  chemical  compound  and  not  a  mixture  of  the 
chloride  and  hypochlorite.  For  our  present  purpose, 
however,  it  may  be  considered  as  such  a  mixture,  for 
under  most  circumstances  it  acts  as  if  it  were.  When 
treated  with  sulphuric  acid  or  hydrochloric  acid,  bleach- 
ing powder  gives  up  hypochlorous  acid  first,  and  then 
chlorine.  The  character  of  the  action  will  be  clear  by 
considering  first  the  conduct  of  the  corresponding  potas- 
sium compounds.  When  a  mixture  of  potassium  hypo- 
chlorite and  potassium  chloride  is  treated  with  dilute 
sulphuric  acid,  the  hypochlorite  is  decomposed  with 
liberation  of  hypochlorous  acid  and  formation  of  potas- 
sium sulphate : 

2KC1O  +  H2SO4  =  K2SO4  +  2HC1O. 

At  the  same  time  the  sulphuric  acid  acts  upon  the  chlo- 
ride liberating  hydrochloric  acid  : 

2KC1  +  H2SO4  =  K2SO4  +  2HC1. 

But  when  hydrochloric  acid  and  hypochlorous  acid  are 
brought  together  they  react  as  represented  in  this 
equation : 

HC1O  +  HC1  =  H2O  +  C12. 

So  that  the  result  of  treating  a  mixture  of  hypochlorite 
and  chloride  with  sulphuric  acid  is  the  liberation  of 
chlorine  : 

KC1O  +  KC1  +  H2SO4  =  K2S04  +  H2O  +  C12. 
With  bleaching  powder  the  reaction  is  : 
Ca(01O)2  +  CaCl2  +  2H2SO4  =  2CaSO4  +  2H2O  +  201,. 

Hypochlorous  acid  can  also  be  made  by  passing  chlo- 
rine gas  into  water  in  which  mercury  oxide  is  suspended. 
The  reaction  is  : 

HgO  +  H2O  +  2C12  =  HgCl2  -f  2HOCL 


118  INORGANIC  CHEMISTRY. 

The  concentrated  solution  of  hypochlorous  acid  has  a 
peculiar  odor  suggesting  that  of  chlorine.  It  is  the  odor 
which  is  familiar  as  that  of  bleaching  powder  or  chloride 
of  lime.  The  acid  undergoes  decomposition  very  readily, 
forming  chlorine  and  a  compound  of  chlorine  and  oxygen. 
A  solution  of  the  acid  bleaches  about  as  well  as  chlo- 
rine, and  when  bleaching  powder  is  used  for  bleaching  it 
is  largely  the  hypochlorous  acid  set  free  from  the  hypo- 
chlorite  which  effects  the  desired  changes. 

Like  chloric  acid,  hypochlorous  acid  is  an  excellent 
oxidizing  agent,  and  is  used  in  the  laboratory  in  this 
capacity. 

Chlorous  Acid,  HClOa. — Although  a  substance  of  this 
composition  is  not  known,  a  number  of  compounds  have 
been  made  which  are  closely  related  to  it.  Such,  for 
example,  are  the  compounds  potassium  chlorite,  KC1O2, 
silver  chlorite,  AgClO2,  etc.  Potassium  chlorite  is  formed 
when  a  solution  of  chlorine  dioxide,  C1O2,  in  water  is 
treated  with  a  solution  of  potassium  hydroxide. '  From 
the  solution  of  the  potassium  compound,  the  silver  com- 
pound can  be  made  by  adding  a  solution  of  silver  nitrate. 

Perchloric  Acid,  HC1O4. — When  the  preparation  of 
oxygen  by  heating  potassium  chlorate  was  considered, 
it  was  pointed  out  that  in  the  first  stage  of  the  decompo- 
sition a  reaction  of  this  kind  takes  place  : 

2KC103  =  KC1  +  KC104  +  02. 

The  compound  KC1O4,  or  potassium  perchlorate,  can  be 
separated  from  the  chloride  by  treating  the  mixture  with 
cold  water  in  which  the  chloride  is  easily  soluble,  while 
the  perchlorate  is  practically  insoluble.  From  the  per- 
chlorate, perchloric  acid  can  be  made  in  the  same  way 
that  chloric  acid  is  made  from  potassium  chlorate,  by 
treating  with  fluosilicic  acid  (see  preparation  of  chloric 
acid).  Perchloric  acid  is,  however,  much  more  stable  in 
concentrated  solution  than  the  other  oxygen  compounds 
of  chlorine,  and  if  the  perchlorate  is  treated  with  sul- 
phuric acid,  perchloric  acid  can  be  obtained  from  the 
mixture  by  distillation.  .  »  ;  i  > 


COMPOUNDS  0V  CHLORINE— GENERAL.  119 

Pure  perchloric  acid,  HC1O4,  can  be  obtained  in  the 
form  of  a  colorless  fuming  liquid.  It  is  a  dangerous  sub- 
stance to  deal  with,  as  it  produces  bad  wounds  when 
brought  in  contact  with  the  flesh,  and  is  very  unstable  and 
explosive.  In  contact  with  combustible  substances  in  gen- 
eral  it  causes  explosion  in  consequence  of  the  ease  with 
which  it  gives  up  oxygen  and  converts  the  combustible 
substances  into  gaseous  products. 

A  hydrate  of  the  formula  HC1O4  +  H2O  or  H3C1O6 
is  known.  Further,  there  are  some  facts  known  that 
point  to  the  existence  of  a  second  hydrate  of  the  formula 
HC104  +  2H,0. 

General. — From  the  above  it  will  be  seen  that  the  com- 
pounds of  chlorine  with  hydrogen  and  oxygen  form  a 
series,  the  members  of  which  bear  a  simple  relation  to 
one  another.  Beginning  with  hydrochloric  acid  the  series 
is  as  follows : 

Hydrochloric  acid,      ....  HC1 

Hypochlorous  acid,    ....  HC1O 

Chlorous  acid, HC1O3 

Chloric  aid, HC1O3 

Perchloric  acid,      .     .     .     .     .  HC1O4 

The  successive  members  differ  from  each  other  by  one 
atom  of  oxgen,  the  ratio  between  the  hydrogen  and  the 
chlorine  remaining  the  same  throughout  the  series. 
While  the  compounds  differ  markedly  from  one  another 
in  many  ways,  they  have  some  common  features.  Upon 
metals,  and  their  oxides  and  hydroxides,  all  the  members 
of  the  series  act  in  general  in  the  same  way  that  hydro- 
chloric acid  does,  the  result  being  the  formation  of 
products  which  do  not  contain  hydrogen,  but  do  contain 
a  metal  in  the  place  of  the  hydrogen.  We  have  examples 
of  these  compounds  in  potassium  chlorate,  KC1O3,  cal- 
cium hypochlorite,  Ca(ClO)2,  potassium  chlorite,  KC1O2, 
and  potassium  perchlorate,  KC1O4.  All  these  compounds 
belong  to  the  class  called  salts,  which  will  presently  be 
taken  up.  On  the  other  hand,  while  there  is  a  class  of 
elements  upon  which  hydrochloric  acid  does  not  act,  the 


120  INORGANIC  CHEMISTRY. 

oxygen  compounds  of  the  above  series  will  in  many 
cases  act  upon  these  elements  and  convert  them  into 
oxides.  Thus  sulphur  and  phosphorus,  which  are  not 
acted  on  by  hydrochloric  acid,  are  converted  into  oxides 
by  the  oxygen  compounds  of  chlorine. 

Finally,  the  addition  of  oxygen  to  hydrogen  and  chlo- 
rine decreases  the  stability  of  the  compound.  Hydro- 
chloric acid,  for  example,  is  characterized  by  great 
stability,  while  hypochlorous  acid,  HC1O,  as  well  as  all 
the  other  members  of  the  series,  is  characterized  by 
instability.  The  larger  the  proportion  of  oxygen,  how- 
ever, the  greater  the  stability  of  the  compound.  The 
most  stable  member  of  the  series  of  oxygen  compounds 
is  perchloric  acid.  Another  fact  that  is  worthy  of  special 
notice  is  that  the  metal  derivatives  or  salts  of  these  acids 
are  more  stable  than  the  acids  themselves.  Many  of  them 
can  be  heated  to  a  comparatively  high  temperature  without 
undergoing  decomposition.  This  is  most  marked  in  the 
case  of  the  perchlorates.  It  will  be  remembered  that  in 
decomposing  potassium  chlorate  for  the  purpose  of  mak- 
ing oxygen  the  change  takes  place  in  two  stages.  In  the 
first,  potassium  perchlorate  is  formed.  In  order  to  de- 
compose this,  however,  the  temperature  must  be  raised 
considerably  higher  than  that  which  was  required  to  effect 
the  breaking  down  of  the  chlorate. 

Compounds  of  Chlorine  with  Oxygen. — The  compounds 
of  chlorine  with  oxygen  are : 

Chlorine  monoxide,  C12O,  and  chlorine  dioxide,  C1Q2. 

The  first  or  chlorine  monoxide,  C12O,  is  formed  by  the 
action  of  chlorine  on  dry  mercury  oxide  : 

HgO  +  2C12  =  HgCla  +  CltO. 

It  is  a  gas  which  can  easily  be  condensed  to  the  liquid 
form.  The  specific  gravity  of  its  vapor  gives  the  molec- 
ular weight  corresponding  to  the  formula  C12CX  It  is 
extremely  unstable,  breaking  down  under  the  influence 
of  heat  into  chlorine  and  oxygen.  With  water  it  forms 
hypochlorous  acid,  thus : 


COMPOUNDS  OF  CHLORINE  WITH  OXYGEN.        121 

ClaO+HaO  =  2HOCl. 

A  substance  formed  by  treating  a  mixture  of  arsenic 
trioxide,  As2O3,  and  potassium  chlorate  with  nitric  acid 
and  by  other  methods  has  been  described  as  a  greenish- 
yellow  gas  which  can  be  condensed  to  an  extremely  un- 
stable liquid  ;  and  it  is  generally  referred  to  under  the 
name  chlorine  trioxide,  the  formula  C12O3  being  ascribed 
to  it.  The  most  careful  investigation  of  this  substance 
has,  however,  shown  that  it  is  not  chlorine  trioxide,  but 
a  mixture  of  chlorine  dioxide  with  varying  quantities  of 
chlorine  or  free  oxygen. 

Chlorine  dioxide,  C1O2,  is  a  greenish-yellow  gas  of  great 
instability.  It  can  be  condensed  to  a  liquid  which  boils 
at  +  9°.  It  is  always  one  of  the  products  of  the  action 
of  concentrated  sulphuric  acid  upon  potassium  chlorate, 
and  is  formed  in  consequence  of  the  decomposition  of 
the  chloric  acid  which  is  first  set  free  : 

2KC1O3  +  H2S04   =  K2S04  +  2HC1O,  ; 
6HC103  =  2HC104  +  2H3O  +  4ClOa 


a. 


When  moderately  dilute  hydrochloric  acid  acts  upon 
potassium  chlorate  a  greenish-yellow  gas  is  formed 
which  has  been  called  euchlorine.  It  is  a  mixture  of 
chlorine  and  chlorine  dioxide,  formed  thus  : 

2KC103  +  4HC1  =  2KC1  +  2H2O  +  201Oa  +  C12. 

Combustible  substances  burn  in  chlorine  dioxide  with 
violence.  This  action  can  be  shown  by  putting  a  few 
small  pieces  of  phosphorus  under  water  in  a  glass  vessel, 
and  upon  this  a  little  potassium  chlorate.  If  now  a  few 
drops  of  concentrated  sulphuric  acid  are  added  through  a 
long  narrow  tube  or  pipette,  the  phosphorus  will  be  seen 
to  burn  under  water.  This  is  due  to  the  liberation  of  chlo- 
rine dioxide,  and  the  action  of  this  compound  upon  the 
phosphorus. 

When  a  solution  of  chlorine  dioxide  in  water  is  treated 
with  potassium  hydroxide,  potassium  chlorite,  KC1O2, 
is  formed  (see  p.  118). 


122  INORGANIC  CHEMISTRY. 

Constitution  of  the  Compounds  of  Chlorine  with  Hydro* 
gen  and  Oxygen. — To  determine  the  constitution  of  un- 
stable compounds  which,  when  they  break  down  at  all, 
are  almost  completely  disintegrated  is  difficult,  and  in 
many  cases  impossible.  Our  knowledge  of  the  constitu- 
tion of  the  oxygen  acids  of  chlorine  is  for  this  reason 
extremely  limited.  Regarding  the  series  of  these  com- 
pounds and  comparing  their  composition  with  that  of 
hydrochloric  acid,  it  would  appear  that  they  are  com- 
pounds of  hydrochloric  acid  with  different  amounts  of 
oxygen.  But  we  have  seen  that  chlorine  and  hydrogen 
are  univalent  elements,  as  is  shown  in  hydrochloric  acid, 
H-C1.  If  each  of  the  atoms  in  the  molecule  of  hydro- 
chloric acid  is  doing  all  it  can  in  holding  the  other  atom 
in  combination,  then,  plainly,  it  is  impossible  for  the 
molecule  to  take  up  oxygen  directly  and  form  a  com- 
pound of  the  formula  H-C1-O  or  O-H-C1.  On  the 
other  hand,  it  is  possible  to  conceive  of  the  atoms  chlo- 
rine, hydrogen,  and  oxygen  as  being  unite'd  in  such  a 
way  that  hydrogen  and  chlorine  shall  be  univalent  and 
oxygen  bivalent.  This  is  represented  in  the  formula 
Cl— O— H.  Further,  extensive  study  of  compounds  con- 
taining hydrogen  and  oxygen  has  made  it  appear  ex- 
tremely probable  that  in  them  the  hydrogen  is  generally 
in  combination  with  oxygen  as  represented  in  the  above 
formula,  and  as  is  represented  also  in  the  formulas 
of  such  compounds  as  potassium  hydroxide,  K— O-H, 

O— H 

calcium    hydroxide,    Ca  <  QTT  >  aluminium    hydroxide, 

/O-H 

Al^-O-H  ,  etc.     All  the  reasons  for  this  cannot  possibly 
X0-H 

be  made  clear  without  a  knowledge  of  a  great  many  facts 
which  must  be  acquired  gradually.  As  regards  such 
hydroxides  as  those  just  referred  to,  the  views  expressed 
in  the  formulas  given  are  certainly  simpler  than  any 
others  which  have  been  proposed,  and  they  are  not  con- 
tradicted by  any  known  facts.  Assuming  then  that  in 
the  acids  of  chlorine  the  hydrogen  is  in  combination 
with  oxygen,  or  that  the  compounds  are  hydroxides,  we 


CONSTITUTION  OF  COMPOUNDS  OF  CHLORINE.      123 

have  the  formulas  H-O-C1,  H-O-C1O,  H-O-ClOa,  and 
H-O-C1O3  for  the  four  compounds.  If,  however,  chlo- 
rine is  univalent  the  additional  oxygen  cannot  be  in 
direct  combination  with  the  chlorine,  and  the  only 
way  in  which  the  constitution  of  these  compounds  can 
be  represented  on  the  assumption  that  hydrogen  and 
chlorine  are  univalent  and  oxygen  bivalent  is  this: 
H-O-C1,  H-O-O-C1,  H-O-O-O-C1,  and  H-O-O-O-O-C1. 
These  formulas  have  been  used  for  some  time,  but 
strong  opposition  has  been  raised  to  them  and  they  are 
now  rapidly  losing  ground.  They  represent  compounds 
in  which  oxygen  atoms  are  in  combination  with  oxygen 
atoms.  But  judging  by  the  conduct  of  hydrogen  dioxide 
and  ozone,  this  kind  of  combination  is  a  very  unstable 
one.  The  acids  of  chlorine  are  unstable  enough,  but  the 
one  which  contains  the  most  oxygen  is  the  most  stable, 
and  this  we  should  hardly  expect  if  the  oxygen  atoms 
are  arranged  as  shown  in  the  above  formulas.  This  is 
not  a  fatal  objection  to  these  formulas,  but  it  makes 
them,  at  least,  appear  improbable. 

The  view  which  now  finds  most  support  is  based  upon 
the  conception  that  the  valence  of  chlorine  towards  oxy- 
gen and  towards  oxygen  and  hydrogen,  or  towards  hy- 
droxyl  as  the  group  O-H  is  called,  varies  from  univa- 
lence  to  septivalence  ;  that  it  is  univalent  in  hydrochloric 
acid  and  in  hypochlorous  acid,  trivalent  in  chlorous 
acid,  quinquivalent  in  chloric  acid,  and  septivalent  in 
perchloric  acid.  This  is  shown  in  the  formulas 

O  O 

H-O-C1,  H-O-C1=O,  H-O-C1  and  H-O-C1=O. 


From  a  study  of  other  similar  compounds  these  com- 
pounds are  regarded  as  derived  from  hydroxides,  thus  : 
Chlorous  acid,O=Cl-O-H,  is  supposed  to  be  formed  from 


the  hydroxide  Cl^—  OH   by  loss  of   water  ;    chloric  acid 
\OH 


124  INORGANIC  CHEMISTRY. 

from  the  hydroxide  C1(OH)5  by  loss  of  two  molecules  of 
water  : 

C1(OH)5  =  C1O2(OH)  +  2H2O  ; 

perchloric  acid  from  the  hydroxide  C1(OH)7  by  the  loss 
of  three  molecules  of  water  : 


C1(OH)7  =  ClO3(pH)  +  3H3O. 

When  the  acids  are  dissolved  in  water  it  is  quite  prob- 
able that  in  many  cases  the  hydroxides  are  formed,  but 
being  unstable  they  cannot  generally  be  isolated.  The 
hydrate  of  perchloric  acid,  HC1O4  +  H,O(H3C1O6),  ap- 
pears to  be  such  a  compound.  It  probably  has  the  con- 
stitution represented  by  the  formula  O2C1(OH)3.  From 
the  compound  C1(OH)7  a  series  of  compounds  can  be 
derived  by  successive  losses  of  one  molecule  of  water,  as 
here  shown  : 

C1(OH)7       -  H90  =  OC1(OH)5  ; 
OC1(OH)6     -  H,0  =  02C1(OH)3;  and 
OaCl(OH)3  -  HaO  =  03C1(OH). 

The  second  and  last  products  are  the  hydrate  and  the 
compound  known  as  perchloric  acid.  While  the  evidence 
in  favor  of  this  view  presented  by  these  compounds 
themselves  is  very  slight,  the  view  is  strongly  supported 
by  the  conduct  of  certain  analogous  derivatives  of  the 
element  iodine,  which  in  many  respects  conducts  itself 
like  chlorine. 

Comparison  of  Chlorine  and  Oxygen.  —  The  power  of 
chlorine  to  combine  with  other  elements  is  nearly  as 
great  as  that  of  oxygen.  It  combines  with  all  other  ele- 
ments except  fluorine,  but  does  not  form  quite  as  great  a 
variety  of  compounds  as  oxygen.  Oxygen  combines  with 
several  elements  in  three  or  four  different  proportions, 
but  chlorine  rarely  combines  in  more  than  two  propor- 
tions with  one  element.  While  in  general  chlorine  and 
oxygen  conduct  themselves  in  the  same  way,  there  is  a 


COMPARISON  OF  CHLORINE  AND  OXYGEN.         125 

very  important  difference  between  them,  to  which  atten- 
tion has  already  been  called  indirectly.  There  are  some 
elements  towards  which  oxygen  has  a  stronger  affinity 
than  chlorine  ;  and  there  are  others  towards  which  chlo- 
rine has  a  stronger  affinity  than  oxygen.  The  former  are 
the  non-metals  or  acid-forming  elements ;  the  latter  are 
the  metals  or  base-forming  elements.  The  difference  is 
shown  most  readily  by  the  fact  that  the  chlorine  com- 
pounds of  the  acid-forming  elements  are  converted  by 
water  into  oxygen  compounds  or  hydroxides,  while  the 
oxides  or  hydroxides  of  the  base-forming  elements  are 
converted  into  chlorides  by  the  action  of  hydrochloric 
acid.  This  distinction  is  not  a  sharp  one  which  can 
•easily  be  made,  for  the  conduct  of  an  element  is  to  a  con- 
siderable extent  dependent  upon  conditions,  particularly 
of  temperature,  and  a  distinction  which  holds  good  under 
one  set  of  conditions  may  possibly  not  hold  good  under 
another  set.  Still  the  above  statements  in  regard  to  the 
conduct  of  the  metals  and  non-metals  towards  chlorine 
and  oxygen  are  in  accordance  with  well-marked  tenden- 
cies of  the  two  classes  of  elements,  as  will  appear  more 
clearly. 

Chlorine  and  oxygen  being  in  general  similar  elements 
we  should  not  expect  them  to  combine  readily  with  each 
other.  Although  they  do  combine  indirectly  in  a  number 
of  proportions,  none  of  the  compounds  are  stable.  There  is 
a  marked  difference  between  these  compounds  and  hydro- 
chloric acid  as  regards  the  ease  with  which  they  are 
formed,  and  also  as  regards  their  stability.  A  marked 
difference  between  chlorine  and  oxygen  is  also  to  be 
found  in  their  relations  to  hydrogen.  While  one  volume 
-of  chlorine  combines  with  one  of  hydrogen  to  form  two 
volumes  of  hydrochloric  acid  gas,  one  volume  of  oxygen 
combines  with  two  volumes  of  hydrogen,  the  three  vol- 
umes of  gas  condensing  to  two  volumes  of  water  vapor. 
So^also  while,  as  we  say,  one  atom  of  chlorine  combines 
with  one  atom  of  hydrogen,  one  atom  of  oxygen  com- 
bines with  two  atoms  of  hydrogen.  What  the  difference 
between  a  bivalent  and  a  univalent  element  consists  in  is 


126  INORGANIC  CHEMISTRY. 

not  known,  but  that  the  difference  is  something  deep 
seated  appears  from  the  marked  difference  in  conduct 
between  chlorine  and  oxygen  in  combining  with  hy- 
drogen. 


CHAPTER  X. 

ACIDS— BASES— NEUTRALIZATION-SALTS. 

General — One  cannot  deal  with  chemical  phenomena 
without  constant  reference  to  acids,  and  in  the  course  of 
our  study  thus  far  a  number  of  substances  belonging  to 
this  class  have  been  met  with.  It  is  now  time  to  inquire 
what  features  these  substances  have  in  common  which 
lead  chemists  to  call  them  all  acids.  What  is  there  in 
common  between  the  heavy,  oily  liquid,  sulphuric  acid, 
the  colorless  gas,  hydrochloric  acid,  and  the  unstable 
compounds  chloric  and  hydrochlorous  acids?  To  un- 
derstand the  common  features  requires  some  knowl- 
edge of  a  class  of  substances  to  which  attention  has 
already  been  given.  These  are  substances  like  caustic 
potash  and  caustic  soda,  or  potassium  and  sodium  hy- 
droxides which  are  called  alkalies,  which  are  the  most 
marked  representatives  of  the  class  of  substances  known 
as  bases.  These  two  classes,  acids  and  bases,  have 
the  power  to  destroy  the  characteristic  properties  of 
•each  other.  When  an  acid  is  brought  in  contact  with 
a  base  in  proper  proportions,  the  characteristic  prop- 
erties of  both  the  acid  and  the  base  are  destroyed. 
They  are  said  to  neutralize  each  other.  They  form  new 
products  which  are  said  to  be  neutral,  which  means  that 
they  have  not  the  properties  of  an  acid  nor  those  of  a 
base.  This  act  of  neutralization  is  an  extremely  im- 
portant one,  with  which  we  have  constantly  to  deal  in 
chemical  operations. 

A  Study  of  the  Act  of  Neutralization. — The  fact  hav- 
ing been  learned  that  acids  and  bases  neutralize  one 
another,  the  next  thing  to  do  is  to  study  the  act  of  neu- 
tralization as  carefully  as  possible,  and  learn  what  chemi- 
cal changes  are  involved  in  it.  For  this  purpose  we 

(127) 


128  INORGANIC  CHEMISTRY. 

should  select  a  number  of  acids  and  a  number  of  basest 
and  study  their  action  upon  one  another.  We  may  take 
sulphuric,  hydrochloric,  and  nitric  acids  ;  and  potassium,, 
sodium,  and  calcium  hydroxides.  We  know  from  many 
analyses  that  have  been  made  that  the  composition  of 
these  substances  is  as  follows  : 

-  •  + 

Hydrochloric  acid, HC1 

Nitric  acid,  .     .     .    „     .     •     •    -  HNO3 

Sulphuric  acid, H2SO4 

Potassium  hydroxide,     ....  KOH 

Sodium  hydroxide,    .     .     .     .     .  NaOH 

Calcium  hydroxide, Ca(OH)2 

The  first  question  to  be  answered  is  whether,  in  order 
to  effect  neutralization,  definite  quantities  of  the  sub- 
stances are  necessary.  To  decide  this,  solutions  of  the 
acids  and  of  the  bases  should  be  prepared  and  allowed 
to  act  upon  one  another  in  different  proportions.  But 
how  shall  we  determine  whether  the  solutions  we  are- 
working  with  are  acid,  basic,  or  neutral  ?  It  has  been 
found  that  all  acids  have  the  power  to  change  the  color 
of  certain  substances.  For  example,  the  dye  litmus  is 
blue.  If  a  solution  which  is  colored  blue  with  litmus  is 
treated  with  a  drop  or  two  of  an  acid,  the  color  is 
changed  to  red.  If  now  the  red  solution  is  treated  with 
a  few  drops  of  a  solution  of  a  strong  base,  the  blue  color 
is  restored.  There  are  many  other  substances  which 
change  markedly  in  color  by  the  addition  of  acids  or 
bases.  These  facts  furnish  a  means  of  recognizing 
whether  a  solution  is  acid  or  basic.  Now,  suppose  that  to 
a  carefully  measured  quantity  of  one  of  the  acid  solutions 
a  few  drops  of  blue  litmus  is  added.  It  will  at  once 
turn  red.  On  adding  slowly  a  solution  of  one  of  the 
bases  the  color  will  remain  red  as  long  as  the  solution  is 
acid,  but  the  instant  it  is  basic  it  will  turn  blue.  By 
noticing  when  the  change  in  color  takes  place,  it  is  pos- 
sible to  determine  exactly  how  much  of  a  certain  basic 
solution  is  required  to  neutralize  the  quantity  of  the  acid 
solution  taken.  If  it  is  found  in  the  case  studied  that 


NEUTRALIZATION.  129 

to  neutralize  20  cc.  of  the  acid  solution  30  cc.  of  the  basic 
solution  are  required,  then,  using  the  same  solutions,  it 
will  be  found  in  every  experiment  that  the  same  quanti- 
ties are  required  to  effect  neutralization,  or  that  the 
change  of  color  takes  place  whenever  these  proportions 
are  reached.  And  no  matter  how  the  quantity  of  one  of 
the  liquids  is  varied,  the  quantity  of  the  other  required 
for  neutralization  varies  in  the  same  proportion.  A  great 
many  experiments  of  this  kind  have  been  performed  with 
many  different  acids,  and  what  is  true  in  one  case  has 
been  found  true  in  all.  It  appears,  therefore,  that  tfie 
act  of  neutralization  is  a  definite  one,  loJiich  takes  place  be- 
tween definite  quantities  of  acid  and  base;  that  for  a  certain 
quantity  of  base  a  certain  quantity  of  acid  is  required  to 
effect  neutralization,  and  vice  versa. 

The  next  question  to  be  answered  is,  What  is  formed 
when  the  acid  and  base  are  neutralized  ?  To  determine 
this,  larger  quantities  of  acids  should  be  neutralized 
with  bases,  and  the  substance  or  substances-  formed 
should  then  be  studied.  If  hydrochloric  acid  is  neu- 
tralized with  sodium  hydroxide  a  solid  product,  sodium 
chloride,  is  formed.  The  action  takes  place  according 
to  the  following  equation  : 

HC1  +  NaOH  =  NaCl  +  H2O. 

Hydrochloric  acid  and  calcium  hydroxide  act  thus  : 
2HC1  +  Ca(OH)2  =  CaCl2  +  2H2O. 

Nitric  acid  acts  upon  the  three  bases  mentioned  above  as 
represented  in  these  equations  : 


HNO3  +KOH  =KN03  +  H2O  ; 
HNO3  +NaOH  =  NaNO3  +  H2O  ; 
2HN03  +  Ca(OH)2  =  Ca(NO3)2  +  H2O. 

Sulphuric  acid  acts  upon  these  same  bases  thus  : 

H2S04  +2KOH  =  K2S04  +  2H2O  ; 
H2SO4  +2NaOH  =  Na0SO4  +  2H2O  ; 
H,S04  +  Ca(OH)2  =  CaS04  +  2H2O. 


130  INORGANIC  CHEMISTRY. 

The  reactions  which  take  place  in  these  cases  are  typi- 
cal of  all  reactions  between  acids  and  bases.  One  of  the 
products  formed  is  always  water,  the  other  is  a  com- 
pound which  is  without  acid  and  basic  properties,  or 
which  is  neutral  and  differs  from  the  acid  in  that  it 
contains  some  other  element  in  place  of  the  hydrogen. 
This  other  element  is  the  one  which  in  the  base  is  in 
combination  with  hydrogen  and  oxygen  as  a  hydroxide. 
The  simplest  case  is  that  of  hydrochloric  acid  and  either 
potassium  or  sodium  hydroxide : 

HC1  +  KOH  =  KC1  +H20. 

As  has  already  been  stated  (see  p.  Ill),  we  have  here  two 
forces  operating  to  bring  about  the  change  :  (1)  the  ten- 
dency of  hydrogen  to  combine  with  hydroxyl  (OH)  to 
form  water;  and  (2)  the  tendency  of  chlorine  to  unite 
with  potassium.  A  similar  statement  may  be  made 
in  regard  to  every  reaction  between  an  acid  and  a 
base. 

General  Statements. — Considering  the  facts  treated  of 
in  the  last  paragraph,  it  appears  : 

(1)  That  an  acid  contains  hydrogen ; 

(2)  That  a  base  contains  a  metal ; 

(3)  That  when  an  acid  acts  upon  a  base  the  hydrogen 
and  metal  exchange  places  ; 

(4)  That  the  substance  formed  by  substituting  hydro- 
gen for  the  metal  of  the  base  is  water ; 

(5)  That  the  substance  obtained  from  the  acid  by  sub- 
stituting a  metal  for  the  hydrogen  is  neither  an  acid  nor 
a  base,  but  is  generally  neutral. 

The  last  statement  is  subject  to  some  modification,  for 
reasons  which  in  some  cases  are  clear  but  in  others  are 
not  apparent.  It  is  true  that  in  some  cases  after  substi- 
tuting a  metal  for  the  hydrogen  the  substance  has  an 
alkaline  reaction,  and  in  other  cases  an  acid  reaction. 

Definitions. — We  have  already  seen  that  hydrochloric 
acid  and  sulphuric  acid  act  upon  certain  metals,  as  iron 
and  zinc,  and  that  the  action  consists  in  giving  up  hy- 
drogen and  taking  up  metal  in  its  place.  The  products 


REACTION  BETWEEN  ACIDS  AND  BASES.  131 

of  this  action  are  the  same  in  character  as  those  formed 
by  the  action  of  acids  on  bases. 

An  acid  is  a  substance  containing  hydrogen,  which  it 
easily  exchanges  for  a  metal,  when  treated  with  a  metal 
itself,  or  with  a  compound  of  a  metal,  called  a  base. 

A  base  is  a  substance  containing  a  metal  combined 
with  hydrogen  and  oxygen.  It  easily  exchanges  its  metal 
for  hydrogen  when  treated  with  an  acid. 

The  products  of  the  action  of  an  acid  on  a  base  are, 
first,  water,  and,  second,  a  neutral  substance  called  a  salt. 

In  the  examples  above  cited  the  products  KNO3,  po- 
tassium nitrate  ;  NaNO3,  sodium  nitrate ;  Ca(NO3)2,  cal- 
cium nitrate  ;  KQSO4,  potassium  sulphate  ;  Na2SO4,  sodium 
sulphate  ;  CaSO4,  calcium  sulphate,  are  salts.  The  rela- 
tions between  them  and  the  acids  from  which  they  are 
derived  will  be  easily  recognized  on  comparing  their 
formulas  with  those  of  the  acids. 

Comparison  of  the  Reaction  between  Acids  and  Hy- 
droxides, and  between  Acids  and  Chlorides. — The  reac- 
tion between  acids  and  hydroxides,  or,  as  it  is  generally 
spoken  of,  between  acids  and  bases,  is  quite  similar  in 
character  to  that  which  takes  place  between  some  acids 
and  chlorides.  This  is  illustrated  by  the  reaction  be- 
tween sulphuric  acid  and  sodium  chloride,  represented 
by  the  equation 

2NaCl  +  H2SO4  =  NauSO4  +  2HC1. 

Here,  as  when  the  hydroxide  is  used,  the  acid  is  neutral- 
ized and  the  salt,  sodium  sulphate,  Na2SO4,  is  formed. 
The  other  product,  however,  is  hydrochloric  acid  instead 
of  water.  For  the  sake  of  closer  comparison  the  two  reac- 
tions may  be  written  thus  : 

NaOH  ,    H  )  Qn    _  Na  )  Qn    ,  HOH  . 
NaOH  +  H  f  SU<  -  Na  f  M<J<  +  HOH  > 

NaCl     ,     H  )  «0       Na  )  so    ,  HC1 
NaCl  '  -  H  f  ^  =  Na  f  ^'°<  +  HC1. 

The  two  reactions  are  thus  seen  to  be  of  the  same  general 
character.  That  with  the  chloride  does  not  take  place 


132  INORGANIC  CHEMISTRY. 

as  readily  as  that  with  the  hydroxide,  and  therefore  is 
not  as  general.  There  are  many  acids  which  have  not 
the  power  to  decompose  chlorides  as  sulphuric  acid  does  ; 
whereas,  in  general,  any  acid  is  neutralized  by  any  me- 
tallic hydroxide.  In  some  cases  this  reaction  is  an  ener- 
getic one  accompanied  by  a  great  evolution  of  heat ;  in 
others  the  reaction  is  not  at  all  energetic.  Both  acids 
and  bases  differ  very  markedly  from  one  another  in  some 
property  which  is  spoken  of  in  a  vague  sort  .of  way  as 
the  strength.  For  the  present  it  is  sufficient  to  recog- 
nize that  this  difference  is  similar  to  the  difference  no- 
ticed between  elements.  Hydrogen  and  chlorine,  for 
example,  differ  markedly  in  their  power  to  act  upon 
other  substances,  and  chlorine  is  spoken  of  as  the  more 
energetic  or  active  element. 

Other  Similar  Reactions. — There  are  many  other  reac- 
tions like  those  which  take  place  between  acids  and 
chlorides,  and  between  acids  and  hydroxides.  Another 
example  is  furnished  by  the  sulphides  and  hydrosuphides, 
which  are  compounds  that  in  some  respects  resemble 
oxides  and  hydroxides.  The  reactions  which  take  place 
between  the  sulphur  compounds  and  acids,  and  between 
the  oxygen  compounds  and  acids,  are  entirely  analogous, 
as  shown  in  the  following  equations  : 

K2S         +  2HC1   =  2KC1    +  H2S  ; 
K2O         +2HC1    =:2KC1    +H2O; 
CaS         +  H2S04  =  CaS04  +  H2S  ; 
CaO         +  H2S04  =  CaS04  +  H2O  ; 
KSH       +HC1     =KC1      +  H2S; 
KOH      +HC1     =KC1     +H2O; 
Ca(SH)2  +  H2S04  =  CaS04  +  2H2S ; 
Ca(OH)2  +  H2S04  =  CaS04  +  2H2O. 

The  product  formed  in  place  of  water  is  the  correspond- 
ing compound  of  sulphur,  H2S.  It  will  be  observed  that 
the  hydrosulphides,  or  compounds  which  have  the  general 
composition  MSH,  neutralize  the  acids  in  the  same  sense 
that  the  hydroxides  do.  If  hydroxides  were  not  known, 
our  conceptions  of  acids  might  easily  be  based  upon  the 


DISTINCTION  BETWEEN  ACIDS  AND  BASES.       133 

relations  of  compounds  to  the  hydrosulphides,  and  the 
substances  now  classed  with  the  acids  would  be  classed 
with  them  upon  this  basis.  As  we  go  on  we  shall  see  that 
there  are  other  reactions  of  the  same  general  character. 

Distinction  between  Acids  and  Bases. — Although  there 
is  no  difficulty  in  distinguishing  between  most  acids  and 
most  bases,  there  are  some  compounds  which  act  some- 
times in  one  way  and  sometimes  in  the  other.  Sulphuric 
acid,  nitric  acid,  and  hydrochloric  acid  always  act  as 
acids,  and  sodium  and  potassium  hydroxides  always  act 
as  bases,  but  some  substances  which  are  generally  basic 
will  under  some  circumstances  act  as  acids,  and  some 
which  act  as  acids  will  occasionally  act  as  bases.  What 
is  the  standard?  How  shall  we  tell  whether  a  sub- 
stance is  an  acid  or  a  base  ?  We  may  take  a  pronounced 
acid,  such  as  hydrochloric  acid,  and  say  that  any  hy- 
droxide which  has  the  power  to  neutralize  this  acid  and 
form  with  it  a  salt  shall  be  called  a  base ;  and  in  the 
same  way  we  may  take  a  pronounced  base,  like  potas- 
sium hydroxide,  and  say  that  any  hydroxide  which  has 
the  power  to  neutralize  this  shall  be  called  an  acid. 
Having  made  the  division  in  this  way,  it  would  be  found 
that  a  few  substances  would  be  included  in  both  lists,  or, 
in  other  words,  some  substances  which  are  basic  toward 
hydrochloric  acid  are  acid  toward  potassium  hydroxide. 
As  an  example,  we  may  take  aluminium  hydroxide, 
A1(OH)3.  This  neutralizes  hydrochoric  acid  and  forms 
aluminium  chloride  according  to  the  equation 

A1(OH)3  +  3HC1  =  A1G1,  +  3H2O. 

But  it  also  neutralizes  potassium  hydroxide  according  to 
the  equation 

A1(OH)3  +  3KOH  =  A1(OK)3  +  3H2O. 

It  may  be  said  in  regard  to  this  case,  as  in  regard  to 
most  other  cases  of  the  kind,  that  the  hydroxide  in  ques- 
tion is  basic  toward  nearly  all  substances  toward  which 
potassium  hydroxide  is  basic  ;  whereas  it  is  acid  toward 
only  three  or  four  of  the  most  energetic  bases.  Bearing 


134  INORGANIC  CHEMISTRY. 

in  mind,  then,  the  fact  that  there  are  some  exceptional 
cases,  it  may  be  said  that  the  distinction  between  acids 
and  bases  is  easily  recognized. 

Metals  or  Base-forming  Elements. — The  question,  What 
is  a  metal  ?  may  fairly  be  asked.  But  unfortunately  it  is 
by  no  means  an  easy  matter  to  give  a  satisfactory  answer 
to  the  question.  We  can  give  examples  of  metals,  such 
as  iron,  zinc,  silver,  calcium,  magnesium,  etc. ;  but  when 
we  attempt  to  find  the  distinguishing  features  of  these 
substances  we  are  somewhat  at  a  loss  to  state  them.  In 
general,  it  may  be  said  that  to  the  chemist  any  element 
is  a  metal  which  with  hydrogen  and  oxygen  forms  a  base, 
or  a  product  which  has  the  power  to  neutralize  acids. 
In  general,  any  element  which  has  the  power  to  enter 
into  an  acid  in  the  place  of  the  hydrogen  is  called  a 
metal,  or  is  said  to  have  metallic  properties.  This  is  the 
sense  in  which  the  word  metal  is  used  in  this  book.  A 
better,  though  a  longer,  name  for  the  metals  is  base-form- 
ing elements. 

Constitution  of  Acids  and  Bases. — As  has  been  pointed 
out,  the  bases  are  hydroxides,  and  these  hydroxides  are 
regarded  as  derived  from  water  by  the  replacement  of 
the  hydrogen  by  metals.  Examples  of  the  hydroxides 
of  univalent,  bivalent,  and  trivalent  metals  were  given 
in  a  previous  chapter  (see  pp.  83-84).  Similarly,  the 
acids  which  contain  oxygen  are  regarded  as  hydroxides, 
or  as  derived  from  water,  as  was  stated  when  the  sub- 
ject of  the  constitution  of  the  acids  of  chlorine  was  under 
consideration.  This  view  is  illustrated  by  the  following 
formulas  of  some  of  the  more  common  acids  : 

Nitric  acid, (HO)NO, 

Sulphuric  acid, (HO)2SO2 

Phosphoric  acid,      .     .    *  , .-     .  (HO)3PO 

Carbonic  acid,     ....     .     .  (HO)2CO 

Metaphosphoric  acid,  ....  (HO)PO2 

Nitrous  acid, .  (HO)NO 

Arsenious  acid, (HO)3As 

Hypochlorous  acid,      ....  (HO)C1 

Perchloric  acid, (HO)CIO, 


CONSTITUTION  OF  ACIDS  AND  BASES.  135 

There  are  three  classes  of  acids  represented  in  this 
list:  (1)  those  with  one,  (2)  those  with  two,  and  (3) 
those  with  three  atoms  of  hydrogen  in  the  molecule.  Or, 
considering  the  compounds  as  hydroxides,  these  classes 
are  :  (1)  those  derived  from  one  molecule,  (2)  those  de- 
rived from  two  molecules,  and  (3)  those  derived  from  three 
molecules  of  water  by  replacement  of  half  the  hydrogen 
by  something  else.  It  is  interesting  to  observe,  also,  that 
this  something  which  replaces  the  hydrogen  is  in  most 
cases  an  element  in  combination  with  oxygen  or,  if  it  is 
not  in  combination  with  oxygen,  it  has  the  power  to  take 
up  more  oxygen.  Thus  hypochlorous  and  arsenious 
acids  are  regarded  as  derived  from  water  by  the  replace- 
ment of  hydrogen  in  water  by  chlorine  and  arsenic  as 


shown  thus  :  H-O-C1  and  H-O-)As.     But,  in  each  case, 

H-0/ 

the  element  which  is  in  combination  with  hydroxyl  has 
the  power  to  combine  with  oxygen.  Hypochlorous  acid 
forms  the  products  (HO)CIO,  (HO)C1O2,  and  (HO)C1O3, 
while  arsenious  acid  forms  arsenic  acid  (HO)3AsO. 

We  may  consider  water  as  forming  the  connecting  link 
between  the  oxygen  acids  and  bases.  If  A  stands  for  any 
acid-forming  element,  and  B  for  any  base-forming  ele- 
ment, then  the  general  formula  of  a  base  is  B(OH),  and 
that  of  an  oxygen  acid  A(OH)  or  OXA(OH),  in  which  Ox 
stands  for  some  number  of  oxygen  atoms  from  one  to 
three  or  four.  We  should  then  have  these  relations  : 


Water.  Acids. 

I.  B'(OH)  HOH         (OXA)'(OH) 

II.B"(OH),  ggg         (OxAr(OH)2 

HOH 

III.  B///(OH)3         HOH         (OXA)///(OH)3 
HOH 

In  these  general  formulas  B"  means  any  bivalent  metal, 
and   B//7.any  trivalent   metal;   and   (OXA)"  means  any 


136  INORGANIC  CHEMISTRY. 

group  of  atoms  which  has  the  power  to  hold  two  hydroxyl 
groups  in  combination,  and  is  therefore  bivalent  like  (O9S), 
and  (OXA)'"  means  a  trivalent  group  like  (OAs). 

Constitution  of  Salts. — The  view  held  in  regard  to  the 
constitution  of  salts  is  based  directly  upon  those  held 
in  regard  to  the  constitution  of  acids  and  bases.  It 
is  believed  that  when  an  oxygen  acid  acts  upon  a  base 
the  action  takes  place  as  represented  in  the  following 
equation : 


O2N-O-|H  +  H-0|-K  =  O3N-O-K  +  H-O-H ; 
or  in  this  : 

02N-|0-H  +  H|-O-K  =  O2N-O-K  +  H-O-H. 

In  either  case  the  salt  formed  appears  as  the  acid,  the 
hydrogen  of  which  has  been  replaced  by  the  metal. 
"Whether  the  hydroxyl  of  the  base  unites  with  the  hydro- 
gen of  the  acid,  or  the  hydroxyl  of  the  acid  unites  with 
the  hydrogen  of  the  base,  cannot  be  determined ;  and, 
as  far  as  the  constitution  of  the  salt  is  concerned,  it 
evidently  makes  no  difference.  The  case  stands  thus  : 
For  reasons  partly  pointed  out  above,  the  bases  are 
regarded  as  hydroxides  ;  for  similar  reasons  the  acids  are 
also  regarded  as  hydroxides.  Now,  when  an  acid  acts 
upon  a  base  water  and  a  product  which  differs  from  the 
acid  in  having  the  metal  of  the  base  in  place  of  its  hydro- 
gen are  formed.  The  simplest  interpretation  of  this 
reaction  is  that  given  above.  A  case  in  which  there 
appears  to  be  no  room  for  doubt  as  to  what  takes  place 
is  that  of  hydrochloric  acid  and  a  simple  base  like  sodium 
hydroxide : 

Cl-H  +  H-O-Na  =  CINa  +  H-O-H. 

It  is  highly  probable  that  the  reaction  between  acids  and 
bases  is  always  of  this  character. 

Basicity  of  Acids. — In  working  with  acids  and  bases  it 
is  noticed  that  some  acids  have  the  power  to  form  but 


BASICITY  OF  ACIDS.  137 

one  salt  with  a  base  like  potassium  hydroxide,  while 
others  have  the  power  to  form  two  or  more  salts  with 
such  a  base.  Thus,  for  example,  hydrochloric  acid,  HC1, 
and  nitric  acid,  HNO3,  can  form  but  one  salt  with  potas- 
sium hydroxide,  and  the  reactions  are  represented  in  the 
following  equations : 

KOH  +  HC1     =  KC1     +  H2O  ; 
KOH  +  HNO3  =  KN03  +  H2O. 

If  only  half  the  quantity  of  base  which  is  required  to 
neutralize  the  acid  is  added,  half  the  acid  remains  un- 
changed, and  on  evaporating  the  solution  the  excess  of 
acid  will  pass  off.  So  also,  if  only  half  the  quantity  of 
acid  which  is  required  to  neutralize  the  base  is  added, 
half  the  base  will  remain  unchanged.  On  the  other  hand, 
if  an  acid  like  sulphuric  acid  is  taken,  it  is  found  that  this 
has  the  power  to  form  two  distinct  salts  with  potassium 
hydroxide,  in  one  of  which  there  is  twice  as  much  of  the 
metal  as  in  the  other.  The  reactions  are  represented 
thus : 

KOH  +  H2SO4  =  KHSO4  +  H2O  ; 
2KOH  +  H2SO4  =  K2S04    +  H2O. 

If  to  a  given  quantity  of  sulphuric  acid  only  half  the 
quantity  of  potassium  hydroxide  which  is  required  to 
neutralize  it  is  added,  the  first  reaction  takes  place  ;  but 
if  the  act  of  neutralization  is  complete  the  second  reac- 
tion takes  place.  An  acid  of  this  kind  can,  further,  form 
•one  salt  with  two  bases,  in  which  one  of  the  hydrogen 
atoms  of  the  acid  is  replaced  by  one  metal  and  the  other 
by  a  second  metal. 

The  different  properties  of  the  two  kinds  of  acids  re- 
ferred to  are  ascribed  to  differences  in  constitution.  .In 
the  molecule  of  hydrochloric  acid,  as  in  that  of  nitric 
acid,  there  is  but  one  atom  of  hydrogen  according  to 
the  views  at  present  held.  If,  therefore,  the  act  of  neu- 
tralization takes  place  in  each  molecule  it  is  complete, 
and  the  salt  is  said  to  be  a  neutral  or  normal  salt.  In 
sulphuric  acid,  however,  there  are  two  atoms  of  hydrogen 


138  INORGANIC  CHEMISTRY. 

in  each  molecule,  and  either  one  or  both  of  these  may  be 
replaced.  If  only  one  is  replaced,  a  salt  of  the  general 
formula  MHSO4  is  obtained.  This  is  still  an  acid,  while 
also  partly  a  salt.  It  is  in  fact  an  acid  salt  or  a  salt  acid. 

Acids  like  hydrochloric  and  nitric  acids  have  not  the 
power  to  form  acid  salts.  They  are  called  monobasic 
acids.  While  acids  like  sulphuric  acid,  which  can  form 
two  salts  with  one  base,  one  of  which  is  acid,  are  called 
dibasic  acids. 

Monobasic  acids  are  those  which  contain  but  one  re- 
placeable hydrogen  atom  in  the  molecule.  Dibasic  acids 
are  those  which  contain  two  replaceable  hydrogen  atoms 
in  the  molecule. 

Similarly,  there  are  tribasic  acids,  like  phosphoric  acid, 
H3PO4,  arsenic  acid,  H3AsO4,  etc.  ;  tetrdbasic  acids,  like 
pyrophosphoric  acid,  H4P2O7 ;  pentdbasic  acids,  like  peri- 
odic acid,  H5IO6 ;  etc.,  etc.  The  higher  the  basicity  of 
the  acid  the  greater  the  variety  of  salts  it  can  yield. 

Acidity  of  Bases. — Just  as  we  speak  of  monobasic, 
dibasic,  tribasic  acids,  etc.,  so  we  distinguish  between 
bases  of  different  acidity.  Thus  there  are  the  monacid 
bases,  like  potassium  and  sodium  hydroxides,  KOH  and 
NaOH  ;  diacid  bases,  like  calcium  and  barium  hydoxides, 
Ca(OH)2  and  Ba(OH)2 ;  triacid  bases,  like  aluminium  and 
ferric  hydoxides,  Al(OH)3-and  Fe(OH)3 ;  etc.,  etc. 

If  a  monobasic  acid  acts  upon  a  monacid  base,  one 
molecule  of  one  forms  a  salt  with  one  molecule  of  the 
other,  and,  in  general,  no  other  reaction  between  the  two 
is  possicl-j.  If  a  monobasic  acid  acts  upon  a  diacid  base 
two  reactions  are  possible,  just  as  when  a  monacid  base 
acts  upon  a  dibasic  acid.  Thus,  when,  for  example,  hy- 
drochloric acid  acts  upon  zinc  hydroxide,  Zn(OH)2,  two 
reactions  are  possible  : 

Zn(OH)2  +  HC1    =  Zn  <  g^  +  H2O; 
Zn(OH)2  +  2HC1  =  ZnCl2      +  2H2O. 

The  compound  ZnCl(OH)  is  still  basic,  just  as  the  salt 
KHSO4  is  still  acid,  and  it  is  called  a  basic  salt.  Simi- 


SALTS.  139 

larly,  a  triacid  base  can  form  three  salts  with  a  monobasic 
acid  as,  for  example,  in  the  case  of  bismuth  hydroxide 
and  nitric  acid,  in  which  three  reactions  are  possible  : 

(  OH  (  N03 

Bi  J  OH  +  HNO3    =  Bi^  OH  +    H,O; 
OH  OH 


Bij 


OH  ( N03 

OH  +  2HN03  =  -Bi\  N03  +  2H2O; 

OH  (OH 

OH  ( N03 

,  OH  +  3HN03  =  Bi-|  N03  +  3H2O. 
(  OH  (  N03 


The  salts  Bi  j  fojj\  and  Bi  j  Q^2  are  basic  salts  or 

basic  nitrates  of  bismuth,  while  the  salt  Bi(NO3)3  is  the 
neutral  or  normal  salt. 

Salts. — From  the  above  it  appears  that  there  are  three 
classes  of  salts  :  (1)  Normal  salts,  which  are  derived  from 
the  acids  by  replacement  of  all  the  acid  hydrogen  atoms 
by  metal  atoms ;  (2)  Add  salts,  which  are  derived  from 
the  acids  by  replacement  of  part  of  the  hydrogen  by 
metal  atoms ;  and  (3)  Basic  salts,  which  are  derived  from 
the  bases  by  neutralization  of  part  of  the  basic  hydroxyl 
by  acids.  Normal  salts  are  generally  neutral ;  or,  if  by 
a  neutral  substance  is  meant  one  which  has  not  the  power 
to  form  salts  with  acids  nor  with  bases,  then  the  expres- 
sion normal  salt  is  synonymous  with  neutral  salt.  But, 
strange  to  say,  some  normal  salts  have  what  is  called  an 
acid  reaction,  and  others  have  an  alkaline  or  basic  reac- 
tion. Thus  a  normal  salt  of  a  weak  acid  with  a  strong 
base  as  sodium  carbonate,  Na2CO3,  has  an  alkaline  reac- 
tion. So  also  a  normal  salt  of  a  strong  acid  with  a  weak 
base  may  have  an  acid  reaction,  as  in  the  case  of  copper 
sulphate,  CuSO4.  As  generally  used,  the  expression  neu- 
tral salt  means  a  salt  which  exhibits  neither  an  acid  nor 
an  alkaline  reaction. 

In  naming  acid  salts  various  methods  are  adopted.  In 
the  case  of  a  dibasic  acid,  the  only  distinction  necessary 
is  between  the  acid  and  the  normal  salts.  The  expres- 


140  INORGANIC  CHEMISTRY. 

sions  acid  potassium  sulphate  and  normal  potassium  sul- 
phate mean,  of  course,  the  salts  which  have  the  formulas 
KHSO4  and  K2SO4,  and  there  is  no  danger  of  confusion. 
We  may,  however,  use  the  names  mono-potassium  sul- 
phate and  di-potassium  sulphate,  or  primary  and  secondary 
potassium  sulphates.  The  last  names  are  convenient  and 
readily  convey  to  the  mind  the  nature  of  the  salt  spoken 
of.  Just  as  dibasic  acids  yield  primary  and  second- 
ary salts,  so  tribasic  acids  yield  primary,  secondary,  and 
tertiary  salts.  For  example,  phosphoric  acid  yields  three 
classes  of  salts  :  primary  phosphates,  of  the  general  formula 
MH2PO4 ;  secondary  phosphates,  of  the  general  formula 
M2HPO4 ;  and  tertiary  phosphates,  of  the  general  formula 
M3PO4.  The  phosphates  of  the  first  two  classes  are  called, 
in  general,  acid  phosphates.  The  tertiary  phosphate  is 
identical  with  the  normal  phosphate.  In  naming  basic 
salts  there  is  no  difficulty  in  the  simplest  cases.  Thus,  tak- 
ing the  three  bismuth  nitrates  the  formulas  of  which  are 

given  above,  the  one  of  the  formula  Bi  j  /QTT\    is  called 

the  mono-nitrate ;  that  of  the  formula  Bi  •<  XTT  *'a,  the  di- 

nitrate;  and  that  of  the  formula  Bi(NO3)3,  the  tri-nitrate 
or  normal  nitrate. 

There  are  many  cases  which  are  much  more  complicated 
than  any  of  those  referred  to  above.  Thus,  there  are 
basic  salts  formed  by  dibasic  acids  and  diacid  bases,  by 
dibasic  acids  and  triacid  bases,  etc.  There  is,  for  ex- 
ample, a  basic  copper  carbonate  formed  by  the  partial 
neutralization  of  two  molecules  of  copper  hydroxide, 
Cu(OH)2,  by  one  molecule  of  carbonic  acid,  CO(OH)2. 
The  relations  will  be  seen  by  the  aid  of  the  following 
equation,  in  which  the  structural  formulas  of  copper  hy- 
droxide and  of  carbonic  acid  are  used  : 


OH   ,  HO,  ^s\  ^^-004 


OH  Cu<OH 


The  salt  is  basic. 


ACID  PROPERTIES  AND  OXYGEN.  141 

Acid  Properties  and  Oxygen. — Almost  all  those  sub- 
stances which  are  called  acids  contain  oxygen,  as,  for 
example,  nitric  acid,  HNO3 ;  sulphuric  acid,  H2SO4 ;  phos- 
phoric acid,  H3PO4 ;  silicic  acid,  H2SiO3 ;  carbonic  acid, 
H2CO3 ;  boric  acid,  H3BO3 ;  etc.  The  presence  of  oxygen 
in  acids  was  recognized  by  Lavoisier.  As  he  showed 
its  presence  in  acids  to  be  general,  and  as  he  found 
that  several  elements  and  some  compounds  are  con- 
verted into  acids  by  combination  with  oxygen,  he  con- 
cluded that  this  element  is  an  essential  constituent  of  all 
acids,  and  therefore  called  it  oxygen,  a  name  which,  as 
already  stated  (see  p.  28),  means  the  acid-former.  Ac- 
cording to  Lavoisier,  hydrochloric  acid,  like  other  acids, 
contained  oxygen,  and  this  view  prevailed  for  many  years. 
As  has  been  pointed  out  under  the  head  of  Chlorine, 
many  investigations  were  undertaken  with  the  object  of 
determining  whether  this  element  does  or  does  not  con- 
tain oxygen,  the  result  being  to  show  that  in  chlorine, 
and  consequently  in  hydrochloric  acid,  there  is  no  oxygen. 
Several  acids  are  now  known  which  are  like  hydrochloric 
acid  in  this  respect,  but  the  latter  is  the  best  known  ex- 
ample. Similar  compounds  are  hydrobromic  acid,  HBr; 
hydriodic  acid,  HI ;  and  hydrocyanic  acid,  HON.  The 
number  of  these  acids  is,  however,  quite  small,  and  it  is 
undoubtedly  true  that,  of  the  compounds  which  we  com- 
monly call  acids,  by  far  the  larger  number  contain 
oxygen  as  an  essential  constituent.  Further,  some  com- 
pounds which  are  basic  can  be  converted  into  acids  by 
introducing  oxygen  into  them. 

On  the  other  hand,  there  are  many  compounds  which 
do  not  contain  oxygen  which  exhibit  reactions  entirely 
analogous  to  those  of  the  acids.  There  are  for  example 
compounds  containing  sulphur  which  combine  with  sul- 
phides, and  others  containing  chlorine  which  combine 
with  chlorides  in  much  the  same  way  that  the  oxygen 
acids  combine  with  oxides,  and  the  compounds  formed 
are  analogous  to  ordinary  salts,  only  they  contain 
sulphur  or  chlorine  in  place  of  oxygen.  Thus,  there  is 
a  compound  of  arsenic  and  sulphur  of  the  composition 
H3AsS4,  known  as  sulpharsenic  acid,  which  is  analo- 


14:2  INORGANIC  CHEMISTRY. 

gous  to  the  oxygen  compound  arsenic  acid,  H3AsO4.  When 
arsenic  acid  is  treated  with  potassium  hydroxide,  KOH, 
this  reaction  takes  place  : 

H3As04  +  3KOH  =  K3As04  +  3H2O. 

So,  too,  when  sulpharsenic  acid  is  treated  with  potassium 
hydrosulphide  this  reaction  takes  place  : 

H3AsS4  +  3KSH  =  K3AsS4  +  3H2S. 

As  many  such  sulphur  compounds  are  decomposed  by 
water  yielding  the  corresponding  oxygen  compounds,  and 
as  most  such  reactions  must  be  studied  in  solution  in  water, 
a  good  reason  for  the  fact  that  they  are  not  as  numerous 
as  the  oxygen  acids  will  be  seen. 

Just  as  sulphur  acids  act  upon  sulphur  bases  to  form 
sulphur  salts,  so  there  are  what  may  be  called  chlorine 
acids  which  act  upon  chlorine  bases  to  form  chlorine 
salts.  For  example,  there  is  a  compound,  H2PtCl6,  known 
as  chlorplatinic  acid,  which  with  chlorides  forms  well- 
marked  salts : 

H2PtCl6  +  2KC1  =  K2PtCl6  +  2HC1. 

The  product  formed  in  the  reaction  represented  by  this 
equation  is  known  as  potassium  chlorplatinate.  The 
reaction  is  analogous  to  the  following,  in  which  oxygen 
compounds  take  part : 

H2Pt03  +  2KOH  =  K2Pt03  +  H20. 

In  the  chlorine  compounds  two  atoms  of  the  univalent 
element  chlorine  take  the  place  of  each  atom  of  the  biva- 
lent oxygen.  Many  such  compounds  are  known  ;  but  in 
working  with  them  the  same  difficulty  arises  that  was 
referred  to  above  in  speaking  of  the  sulphur  compounds  ; 
many  of  the  chlorides  which  are  capable  of  forming 
chlorine  salts  are  decomposed  by  water  and  converted 
into  oxygen  acids.  Therefore,  if  we  start  with  a  chlorine 
acid  and  work  in  water  solution  the  probability  is  that 


NOMENCLATURE  OF  ACIDS.  143 

the  product  obtained  will  be  an  oxygen  compound.  The 
fact  that  the  oxygen  acids  are  the  most  prominent  is 
partly  to  be  ascribed  to  the  fact  that  water  is  in  such 
general  use  as  a  solvent.  The  analogous  solvent  for 
the  sulphur  compounds  would  be  liquid  hydrogen  sul- 
phide, H2S,  but  at  ordinary  temperatures  this  is  a  gas, 
and  it  is,  therefore,  impossible  to  work  with  the  sulphur 
compounds  under  conditions  analogous  to  those  under 
which  we  work  with  the  oxygen  compounds.  The  same 
statement  applies  to  the  chlorine  compounds  for  which 
the  analogous  solvent  would  be  liquid  hydrochloric  acid, 
HC1, — not  the  solution  of  the  gas  in  water. 

Nomenclature  of  Acids. — The  names  of  the  acids  of 
chlorine  illustrate  some  of  the  principles  of  nomenclature 
in  use  in  chemistry.  The  acid  of  the  series  which  is  best 
known  is  called  chloric  acid.  In  naming  acids  the  suffix 
ic  is  always  used  in  naming  the  principal  member  of  a 
group  of  acids  containing  the  same  elements.  This  is 
seen  in  the  names  hydrochloric,  sulphuric,  nitric,  phos- 
phoric, silicic,  carbonic,  acetic,  etc.  If  there  are  two 
acids  containing  the  same  elements,  that  one  of  the  two 
which  contains  the  smaller  proportion  of  oxygen  is  given 
a  name  ending  in  ous.  Thus  we  have  the  two  series : 

Chloric  acid,     .  .  HC1O3  Chlorous  acid,  .     .  HC1O2 

Sulphuric  acid,  .  H2SO4  Sulphurous  acid,  .  H2SO3 

Nitric  acid,  .     .  .  HNO3  Nitrous  acid,     .     .  HNO2 

Phosphoric  acid,  .  H3PO4  Phosphorous  acid,  H3PO3 

For  most  cases  which  present  themselves  this  method 
of  naming  will  suffice,  but  in  others  the  number  of  acids 
known  is  larger  than  two,  as,  for  example,  in  the  series 
of  chlorine  acids.  In  such  cases  recourse  is  had  to 
prefixes.  If  there  is  an  acid  known  containing  a  smaller 
proportion  of  oxygen  than  the  one  whose  name  ends  in 
ous,  it  is  generally  designated  by  means  of  the  prefix 
hypo,  which  is  derived  from  the  Greek  vno,  signifying 
tinder.  Thus  there  are  the  following  examples' :  Hy- 
pochlorous  acid,  HC1O  ;  hyposulphurous  acid,  H2SOQ ; 
hyponitrous  acid,  H2N2O2;  and  hypophosphorous  acid, 


144  INORGANIC  CHEMISTRY. 

H3PO2.  It  will  be  seen  on  comparing  the  formulas  of 
these  acids  with  those  above  given  that  they  differ  from 
them  in  a  very  simple  way. 

In  the  series  of  chlorine  acids  there  is  one  which  con- 
tains a  larger  proportion  of  oxygen  than  chloric  acid.  It 
is  called  perchloric  acid,  the  Latin  prefix  per  signifying 
here  very  or  fully.  Similarly  there  is  a  perbromic  acid 
and  a  permanganic -acid.  Other  cases  arise,  but  they  are 
of  a  more  or  less  special  character,  and  the  compounds 
are  given  special  names  according  to  circumstances. 

Nomenclature  of  Bases. — As  pointed  out  above,  a  base 
is  a  compound  of  a  metal  with  hydrogen  and  oxygen. 
The  bases  are  commonly  known  as  hydroxides ;  and  in 
order  to  distinguish  between  the  hydroxides  of  the  differ- 
ent metals,  the  names  of  the  metals  are  put  before  the 
name  hydroxide,  as  in  naming  the  oxides  and  chlorides. 
Thus,  as  has  been  seen,  caustic  soda,  NaOH,  is  called 
sodium  hydroxide,  etc.  It  is  necessary  in  some  cases  to- 
distinguish  between  two  hydroxides  of  the  same  metal. 
This  is  done  by  using  the  suffixes  ous  and  ic  in  the  same 
sense  as  they  are  used  in  naming  oxides  and  chlorides. 
Thus  ferric  hydroxide  has  the  composition  Fe(OH)3,  and 
ferrous  hydroxide  the  composition  Fe(OH)2 ;  cuprous  hy- 
droxide is  Cu(OH),  and  cupric  hydroxide  Cu(OH)2,  etc. 
These  compounds  are  sometimes  called  hydrates,  and 
there  are  some  good  reasons  for  using  this  name,  as 
will  be  more  fully  shown  in  a  later  paragraph.  On  the 
other  hand,  compounds  in  which  water  as  such  is  re*- 
garded  as  present  are  called  hydrates,  and  there  is- 
danger  of  confusion  if  the  same  name  is  used  to  desig- 
nate what  are  believed  to  be  two  entirely  different  classes 
of  compounds.  As  examples  of  hydrates  we  have  salts 
with  their  water  of  crystallization,  chlorine  hydrate, 
C19  +  8H2O  ;  hydrochloric  acid  hydrate,  HC1  +  2H2O  ; 
etc.  While  some  of  the  compounds  which  are  commonly 
regarded  as  hydrates  should  probably  be  classed  with  the 
hydroxides,  there  seem  to  be  two  classes,  and  it  is  there- 
fore desirable  to  have  two  names. 

Nomenclature  of  Salts. — Theoretically  every  metal  caji 
yield  a  salt  with  every  acid.  The  salts  derived  from  a 


NOMENCLATURE  OF  SALTS.  145 

given  acid  receive  a  general  name,  and  this  general 
name  is  qualified  in  each  case  by  the  name  of  the  metal 
contained  in  the  salt.  Thus,  all  the  salts  derived  from 
nitric  acid  are  called  nitrates  ;  all  the  salts  derived  from 
chloric  acid  are  called  chlorates  ;  the  salts  of  sulphuric 
acid  are  called  sulphates  ;  *  the  salts  of  phosphoric  acid 
are  called  phosphates;  *  etc.  So  too,  further,  the  salts  of 
chlorous  acid  are  called  chlorites;  those  of  nitrous  acid, 
nitrites ;  those  of  sulphurous  acid,  sulphites ;  etc.,  etc. 
It  will  be  noticed  that  the  final  syllable  of  the  name 
of  the  salt  differs  according  to  the  name  of  the  acid.  If 
the  name  of  the  acid  ends  in  ic,  the  name  of  the  salt  de- 
rived from  it  ends  in  ate.  If  the  name  of  the  acid  ends 
in  ous,  the  name  of  the  salt  ends  in  ite.  To  dis- 
tinguish between  the  different  salts  of  the  same  acid, 
the  name  of  the  metal  contained  in  it  is  prefixed. 
Thus,  the  potassium  salt  of  nitric  acid  is  called  po- 
tassium nitrate,  the  sodium  salt  is  called  sodium  ni- 
trate ;  the  calcium  salt  of  sulphuric  acid  is  called  calcium 
sulphate  ;  the  magnesium  salt  of  nitrous  acid  is  magne- 
sium nitrite ;  the  calcium  salt  of  hypochlorous  acid  is 
calcium  hypochlorite  ;  etc.,  etc.  If  a  metal  forms  two 
salts  with  the  same  acid  in  one  of  which  the  valence  of 
the  metal  is  lower  than  in  the  other,  the  one  in  which 
the  valence  of  the  metal  is  lower  is  designated  by 
means  of  the  suffix  ous,  while  the  one  in  which  the 
valence  of  the  metal  is  higher  is  designated  by  means  of 
the  suffix  ic.  Thus  there  are  two  series  of  salts  of  iron 
which  correspond  to  the  two  chlorides  FeCl2  and  Fed,. 
In  one  series  the  iron  appears  to  be  bivalent,  in  the 
other  trivalent.  Examples  are,  Fe(NO3)2  and  Fe(NO3)3 ; 
FeSO4  and  Fe2(SO4)8 ;  etc.  Those  salts  in  which  the 
iron  is  bivalent  are  called  ferrous  salts,  as  ferrous 
nitrate,  ferrous  sulphate,  etc.  ;  and  those  in  which  it  is 
trivalent  are  called  ferric  salts,  as  ferric  nitrate,  ferric 
sulphate,  etc.  Similarly  there  are  two  series  of  copper 


*  Strictly  speaking,  the  salts  of  sulphuric  acid  should  be  called  sul- 
phurates, and  those  of  phosphoric  acid  phosphorates,  but  for  the  sake  of 
euphony  and  convenience  these  names  are  shortened  to  the  above  forms. 


14:6  INORGANIC  CHEMISTRY. 

salts  known  as  cuprous  and  cupric  salts  ;  and  two  series 
of  mercury  salts  known  as  mercurous  and  mercuric  salts. 
If  the  salts  of  hydrochloric  acid  were  named  in  ac- 
cordance with  the  principle  just  explained,  they  would 
be  called  hydrocMorates,  and  this  name  is  sometimes  used 
for  complex  salts,  but  in  the  case  of  the  salts  of  the 
metals  it  will  be  observed  that  these  are  identical  with 
the  products  formed  by  direct  combination  of  the  metals 
with  chlorine.  Thus,  hydrochloric  acid  and  zinc  act  as 
represented  in  the  equation 

Zn  +  2HC1  =  ZnCl,  +  H2 ; 
while  zinc  and  chlorine  act  thus : 
Zn  +  01,  =  ZnCl,. 

In  each  case  the  same  product,  ZnCl2,  is  formed.  But 
these  compounds  of  metals  with  chlorine  are  called  chlo- 
rides, as  has  already  been  explained.  Hence  for  these 
cases  the  name  hydrocJdorate  is  unnecessary. 

The  name  hydrate  to  which  reference  was  made  in  a 
paragraph  above  suggests  a  salt  of  hydric  acid.  Potas- 
sium hydrate  signifies  the  potassium  salt  of  this  acid  or 
of  water.  In  one  sense  this  is  a  proper  name  for  the 
compound.  It  is  water  in  which  a  part  of  the  hydrogen 
is  replaced  by  a  metal,  and  it  is  in  this  respect  like  a 
salt.  While,  however,  there  is  an  unmistakable  analogy 
between  the  formation  of  a  metallic  hydroxide  from 
water  and  that  of  a  salt  from  an  acid,  it  appears,  on  the 
whole,  wise  not  to  class  water  with  the  acids  nor  with 
the  bases,  but  rather  to  regard  it  as  the  connecting  link 
between  the  two  classes.  We  shall  see  later  that  the 
similar  compounds  hydrogen  sulphide,  H2S,  and  hy- 
drogen selenide,  H2Se,  have  much  more  marked  acid 
properties  than  water.  When  treated  with  metallic  hy- 
droxides they  form  salts  of  the  general  formulas  M,S 
and  MaSe. 


CHAPTER  XI. 

NATURAL  CLASSIFICATION  OF  THE  ELEMENTS— THE 
PERIODIC  LAW. 

Historical. — It  has  long  been  known  that  simple  rela- 
tions exist  between  the  atomic  weights  of  some  elements 
which  resemble  one  another  closely.  Thus  chlorine, 
bromine,  and  iodine  are  very  similar  elements.  Their 
atomic  weights  are  35.18,  79.34,  and  125.89  respectively. 
It  will  be  seen  that  the  atomic  weight  of  bromine,  79.34, 
is  approximately  the  mean  of  those  of  chlorine  and 
iodine.  We  have 

35.18  + 125.89  =  Q()  53 

2 

A  similar  group  is  that  of  sulphur,  selenium,  and  tellu- 
rium, which  resemble  one  another  as  closely  as  chlorine, 
bromine,  and  iodine  do.  The  atomic  weights  are  S  = 
31.83,  Se  =  78.42,  and  Te  =  126.52.  WTe  have  here 

I  *u*  +  i*.n  , 


Other  groups  are  those  of  phosphorus,  30.79,  vanadi- 
um, 50.99,  and  arsenic,  74.44 : 

30.79  + 7444  =  g261; 

lithium,  6.97,  sodium,  22.82,  and  potassium,  38.82 : 
..."        6.97  + 38.82  =  22m 

a 

(147) 


148  INORGANIC  CHEMISTRY. 

In  1863-64  J.  A.  E.  Newlands  called  attention  to  the 
fact  that  if  all  the  elements  are  arranged  in  a  table  in  the 
order  of  their  atomic  weights,  beginning  with  that  one 
which  has  the  lowest  atomic  weight  and  ending  with  that 
one  which  has  the  highest  atomic  weight,  provided  they 
are  arranged  horizontally  in  groups  of  seven,  placing  the 
eighth  under  the  first,  the  ninth  under  the  second,  etc., 
then  similar  elements  would  fall  in  the  same  perpen- 
dicular line.  Newlands'  arrangement  was  quite  imper- 
fect, and  it  required  considerable  modification  in  order  to 
make  it  appear  at  all  satisfactory.  In  1869  and  1870  two 
papers  appeared,  one  by  D.  Mendeleeff  and  the  other  by 
Lothar  Meyer,  in  which  these  relations  are  treated  in  a 
masterly  manner,  and  it  was  then  seen  that  one  of  the 
most  important  laws  of  chemistry  had  been  discovered. 
Everything  learned  since  then  has  only  made  it  appear 
more  and  more  certain  that  the  law  which  is  known  as 
the  periodic  laiv  is  a  fundamental  law  of  chemistry. 

Arrangement  of  the  Elements. — Mendeleeff  and  Lothar 
Meyer  have  proposed  several  arrangements  for  the  pur- 
pose of  making  clear  the  connection  between  the  proper- 
ties and  atomic  weights  of  the  elements.  Those  which 
have  proved  most  useful  will  first  be  given,  and  then  the 
connection  between  the  atomic  weights  and  properties 
will  be  discussed  briefly.  The  different  arrangements 
are  to  be  regarded  only  as  different  ways  of  expressing 
the  same  law,  and  no  one  of  them  is  perfect.  The  inves- 
tigation of  the  relations  between  the  atomic  weights  and 
the  properties  of  the  elements  has  not  yet  been  pushed 
far  enough  to  justify  a  final  opinion  as  to  the  character 
of  the  relations,  but  it  has  nevertheless  reached  a  stage 
in  which  we  are  justified  in  stating  that  these  relations 
are  general  and  deep-seated. 


MENDELEEFF'S  TABLES  OF  THE  ELEMENTS.      149 


TO 

oo'oi 

QO't^ 

WTO 

O  O 

og 

. 

tJ 

"I 

1 

fc    1    £ 

S* 

SJf 

1 

1 

§ 

s's 

§g 

OS  (JJ 

o 

II  II 

II  II 

1 

1 

£5 

S2 

Ao 

. 

i  "^ 

SB  ^~ 

OS 

1 

I 

**•         ^ 

P,    W    o 

1  *  * 

O 

II 
fe 

"   :§ 

S     ii 

0 

ii 

ii 
i 

II 

H 

1 

1 

1 

M 

8 

g 

5 

1 

1 

II        TO 

MM 

ila 

«o 
II 

03      g 

II 

OJ        «0 
02      °» 

II 

n 

1 

II 

o 

0 

0 

i 

i 

TO 

B 

TO 

§ 

I  rf  4 

§    «    « 
O 

II 

II 
P-l       ** 

2 

ll 

II  c, 

3  a6 

II 

g 

m 
II      ^ 

02       II 

5 

oo 

S 

II 

II 

S 

1 

oo 

0» 

00 

Oi 

• 

1 

HH 

II 

II    ^ 

II    10- 

S   w   o 

JSS 

II 
0 

II 
H 

O     OT 

ii 

3 

1 

n    1 

H  1 

fe 

Oi 

z> 

,_, 

I  'i 

o 

II 
PQ 

II 

£ 

§ 

II 

J  s 
ii 

TO 
II        § 

i 

II 

.0 

II 

1 

j 

j 

i 

3 

1 

"* 

1  >  § 

c» 
II 

£     II 

ii   3 

a  " 

S  ; 

II 

o 

1 

6 

M 

1 

1 

^ 

TO 

TO 

OS 

^ 

1-4 

It 

i 

1       0* 

x 

« 

&  i  °: 

W 

5z      os 

n  ^ 

II    TO' 

II 

§  '  « 

*" 

o      II 

'So    1 

3^ 

II 

3 

II 
M 

"  3 

Sg 

1 

^     1 

1 

_       ,,, 

TO       •** 

0       CO 

l>       00 

0       0 

S    S 

02 

150 


INOEGANIG  CHEMISTRY. 


MENDELEEFF'S  TABLE  II. 


I 

I. 

II. 

III. 

IV. 

V. 

VI. 

RaO 

I. 

Li  =  7 

K      39 

Rb    85 

Cs   133 

-      - 

-     - 

RO 

II. 

Be  =  9 

Ca    40 

Sr     87 

Ba  137 

—     — 

—     — 

R303 

III. 

B  =  ll 

Sc     44 

Y      89 

La  138 

Yb  173 

—     — 

RO, 

IV. 

(H4C) 

C  =  12 

Ti     48 

Zr     90 

Ce  142 

—     — 

Th  231 

Ra06 

V. 

(H3N; 

N  =  14 

V      51 

Cb    94 

Di   146 

Ta  182 

_     _ 

RO, 

VI. 

(HaO) 

O  =  16 

Cr     52 

Mo    96 

—     — 

W   184 

U    240 

R307 

VII. 

(HF) 

F  =  19 

Mn   55 

_     _ 

—     — 

—     — 

—     — 

RO4 

Fe    56 

Ru  103 

—     — 

Os  192? 

—     — 

VIII. 

Co    58 

Rh  104 

_     _ 

Ir    193 

—      - 

Ni     59 

Pd  106 

—     — 

Pt   195 

—     — 

RaO 

I. 

H  =  l 

Na=23 

Cu    63 

Ag  108 

—     — 

Au  196 

—     — 

RO 

II. 

Mg    24 

Zn    65 

Cd  112 

—     — 

Hg200 

—     _ 

RaO, 

III. 

Al     2? 

Ga    69 

In    113 

—     — 

Tl    204 

—      — 

RO2 

IV. 

(H4R) 

Si      28 

Ge    72 

Sn  118 

—     — 

Pb  206 

—      — 

R.,06 

V. 

(H3R) 

P       31 

As    75 

Sb  120 

—     — 

Bi    209 

—     — 

RO8 

VI. 

(HaR) 

S       32 

Se     79 

Te  125? 

—     — 

—      — 

_     _ 

R,0T 

VII. 

(HR) 

01  35.5 

Br     80 

I      127 

-     - 

-      - 

-      - 

In  the  above  tables  the  approximate  atomic  weights 
are  used  instead  of  those  which  have  been  determined 
and  calculated  with  the  greatest  care.  For  most  pur- 
poses in  the  laboratory  the  approximate  figures  answer 
well  enough,  and  they  are  most  commonly  used.  In 
the  following  table  of  Lothar  Meyer  the  refined  atomic 
weights  (as  calculated  by  Meyer  and  Seubert)  are  used. 
The  difference  between  the  two  sets  of  figures  is  in  most 
cases  very  slight.  The  atomic  weights  adopted  in  this 
book  are  those  calculated  by  F.  W.  Clarke  (see  "  The 
Constants  of  Nature,"  Part  V,  1897). 


MEYERS  TABLE  OF  THE  ELEMENTS. 


151 


S      85 


"  «sf 


£5         „• 


S      s 


s  a 


152  INORGANIC  CHEMISTRY. 

In  Mendeleeff's  Table  I  the  elements  are  arranged  in 
horizontal  lines,  beginning  with  lithium.  When  the 
eighth  element  in  the  order  of  the  increasing  atomic 
weights  is  reached  it  is  found  that  it  is  very  much  like 
lithium.  It  is  sodium.  If  this  is  placed  below  lithium, 
and  the  next  six  elements  in  the  same  horizontal  line, 
when  the  fifteenth  element  is  reached,  it  is  found  like 
the  eighth  to  be  similar  to  lithium.  Up  to  and  includ- 
ing manganese  there  are  twenty-one  elements  excluding 
hydrogen.  These  fall  then  naturally  into  three  series  of 
seven  members  each,  and  placing  these  horizontally, 
those  elements  which  fall  in  the  same  perpendicular 
lines  have  the  same  general  character.  This  is  seen 
most  strikingly  in  Group  I,  in  which  lithium,  sodium, 
and  potassium  fall,  and  in  Group  Y,  in  which  nitrogen, 
phosphorus,  and  vanadium  fall ;  but  there  is  no  difficulty 
in  recognizing  the  similarity  in  the  other  groups.  The 
three  elements  following  manganese,  viz.,  iron,  nickel, 
and  cobalt,  are  very  much  alike,  and  they  certainly  do 
not  belong  in  Groups  I,  II,  and  III,  while  the  next 
element,  copper,  has  some  properties  which  ally  it  to  the 
members  of  Group  I.  The  next  six  elements  fall  in 
Groups  II  to  VII,  and  are  evidently  in  place,  and  the 
six  following  fall  in  Groups  I  to  YI,  and  are  also  in  their 
proper  places,  as  far  as  their  properties  are  concerned. 
After  molybdenum  in  the  sixth  series  comes  a  blank 
which  means  that  there  is  no  element  to  fill  that  place, 
but  that  probably  there  is  one  undiscovered  which  has 
the  atomic  weight  approximately  100,  and  has  properties 
which  are  similar  to  those  of  manganese.  Then  follow 
three  elements  which  resemble  one  another  as  closely  as 
iron,  nickel,  and  cobalt  do.  These  do  not  belong  in 
Groups  I,  II,  and  III,  but  form  a  small  independent 
group.  These  two  groups  of  three  elements  occur  at  the 
end  of  the  fourth  and  sixth  series  respectively.  We 
should  therefore  expect  to  find  a  similar  group  at  the  end 
of  the  eighth  series.  No  such  group  is  known,  however, 
though  at  the  end  of  the  tenth  series,  where  we  should 
look  for  the  next  similar  small  group,  there  are  the  three 
elements  iridium,  platinum,  and  osmium.  The  elements 


ARRANGEMENT  IN  MENDELEEFF'S  SECOND  TABLE.  153 

of  series  2  beginning  with  lithium  and  ending  with  fluo- 
rine differ  in  some  respects  quite  markedly  from  all  the 
other  elements,  as  will  be  seen  when  they  are  taken 
up.  Beginning  with  sodium,  it  will  be  seen  that  there 
are  two  series  of  seven  elements  and  a  short  series  of 
three ;  then  again  two  series  of  seven  and  a  series  of 
three ;  and,  although  the  following  series  are  imperfect, 
it  is  not  difficult  to  recognize  that  the  same  general  ar- 
rangement of  the  elements  holds  good  to  the  end.  A 
series  of  seven  elements  is  called  a  short  period ;  while 
two  short  periods  with  the  accompanying  three  similar 
elements  constitute  what  is  called  a  long  period. 

In  MendeleefFs  Table  II  the  long  periods  are  ar- 
ranged in  perpendicular  lines,  each  long  period  begin- 
ning and  ending  with  a  short  period  and  having  a  side 
group  of  three  elements  in  the  middle.  Thus  in  the 
column  beginning  with  potassium  there  is,  first,  the 
short  period  potassium  to  manganese,  then  the  side 
group  iron,  cobalt,  nickel,  and  then  the  short  period 
copper  to  bromine.  In  this  table  similar  elements  occur 
in  the  same  horizontal  lines.  Thus  in  one  line  there  are 
lithium,  potassium,  rubidium,  and  caesium  ;  in  another 
sulphur,  selenium,  and  tellurium  ;  and  in  another  chlo- 
rine, bromine,  and  iodine. 

The  symbols  at  the  top  of  each  column  in  Table  I 
have  reference  to  the  general  formulas  of  the  compounds 
which  the  elements  in  each  group  form  with  oxygen  and 
with  hydrogen.  Beginning  with  Group  I,  the  general 
formula  of  the  oxygen  compounds  of  the  members  of  this 
group  is  B2O,  in  which  B  represents  any  element  of  that 
group ;  the  general  formula  of  the  oxygen  compounds  of 
the  members  of  Group  II  is  BO  ;  and  so  on.  It  will  be 
observed  that  the  oxygen  compounds  grow  more  and 
more  complex  from  Group  I  to  Group  VII.  Writing 
BO,  BO2,  and  BO3  with  doubled  formulas,  thus :  B2O2, 
B2O0  and  B2O6,  the  series  of  oxygen  compounds  is 
represented  as  below : 

B90,    B202,    B203,    B204,    B206,    B2O6,    B2O7. 


154  INORGANIC  CHEMISTRY. 

As  regards  the  general  formula  of  the  oxygen  com- 
pounds of  the  members  of  Group  VIII,  it  must  be  said 
that  it  does  not  in  general  correspond  to  the  composition 
EO4.  Osmium  and  ruthenium  do,  however,  form  the 
oxides  OsO4  and  EuO4. 

There  is  also  regularity  in  the  composition  of  the  hy- 
drogen compounds.  Beginning  with  Group  VII,  those 
members  which  combine  with  hydrogen  form  compounds 
of  the  general  formula  EH,  as,  for  example,  C1H,  hydro- 
chloric acid  ;  FH,  hydrofluoric  acid ;  etc.  Those  mem- 
bers of  Group  VI  which  combine  with  hydrogen  form 
compounds  of  the  general  formula  EH2,  as,  for  example, 
water,  H2O,  and  hydrogen  sulphide,  H2S.  The  maxi- 
mum power  of  combining  with  hydrogen  is  met  with  in 
Group  IV,  in  which  occur  the  elements  carbon  and  sili- 
con. These  form  the  hydrogen  compounds  CH4  and 
SiH4.  The  members  of  Groups  I,  II,  and  III  do  not 
readily  form  compounds  with  hydrogen.  A  few  are 
known,  but  they  are  quite  unstable. 

The  hydroxides  vary  in  composition  from  the  simple 
form  E(OH)  to  E(OH)7.  While  in  the  first  four  groups 
well-marked  examples  of  the  hydroxides  E(OH),  E(OH)2, 
E(OH)3,  and  E(OH)4  are  found,  in  the  fifth  group  the 
hydroxides  have  not  the  general  formula  E(OH)5,  though 
several  of  them  have  the  formula  OE(OH)3,  as  phos- 
phoric, arsenic,  and  antimonic  acids,  which  are  respec- 
tively OP(OH)3,  OAs(OH)3,  and  OSb(OH)3.  These  may 
be  regarded  as  derived  from  the  hydroxides  of  the  gen- 
eral formula  E(OH)5  by  loss  of  water  : 

K(OH)S  =  OE(OH)3  +  H,0. 

Hydroxyl  derivatives  of  the  members  of  Group  VI 
corresponding  to  the  general  formula  E(OH)6  are  known, 
as,  for  example,  the  so-called  hydrate  of  sulphuric  acid, 
S(OH)6.  The  maximum  hydroxides  of  Group  VII  should 
have  the  general  formula  E(OH)7,  but  those  known  do 
not  correspond  to  this.  The  nearest  approach  to  it  is 
found  in  crystallized  periodic  acid,  H5IO6,  which  may  be 


LOTHAR  MEYER'S    ARRANGEMENT.  155 

regarded  as  derived  from  the  hydroxide  I(OH)7  by  loss 
of  one  molecule  of  water,  thus  : 

I(OH),  =  OI(OH).  +  H,0. 

The  arrangement  of  Lothar  Meyer  is  a  continuous 
one.  The  elements  are  arranged  on  a  spiral  beginning 
with  lithium  and  ending  with  uranium.  The  divisions 
are  such  that  when  the  two  ends  of  the  table  are  brought 
together  on  a  cylinder,  the  line  ending  with  fluorine  will 
join  that  beginning  with  sodium  ;  that  ending  with  nickel 
will  join  that  beginning  with  copper ;  and  so  on.  In  other 
respects  the  arrangement  is  much  like  that  in  Mende- 
leefF s  Table  I.  "What  Mendeleeff  calls  a  Group,  Lothar 
Meyer  calls  a  Natural  Family,  while  those  elements  which 
fall  in  the  same  horizontal  line  are  said  to  form  a  series. 
Now,  each  natural  family  falls  into  two  groups  indicated 
by  the  letters  A  and  B  placed  above.  Thus  the  first 
natural  family  falls  into  Group  A,  consisting  of  lithium, 
sodium,  potassium,  rubidium,  and  caesium,  and  Group  B, 
consisting  of  copper,  silver,  and  gold.  Family  II  falls 
into  Group  A,  consisting  of  glucinum,  magnesium,  cal- 
cium, strontium,  barium,  and  perhaps  erbium,  and  Group 
B,  consisting  of  zinc,  cadmium,  and  mercury  ;  etc.  The 
members  of  each  group  in  a  family  resemble  one  another 
much  more  closely  than  they  resemble  the  members  of 
the  other  group.  * 

The  figures  at  the  bottom  of  the  table  of  Lothar 
Meyer  refer  to  the  valence  of  the  elements  in  each  group. 
Judging  by  the  composition  of  the  oxides  and  hydrox- 
ides, the  valence  increases  from  1  to  8  from  Families  I  to 
VIII.  But  the  valence  of  the  elements  in  each  family 
varies  according  to  conditions.  Thus  the  valence  of  the 
elements  of  Family  IY  is  generally  4,  as  shown  in  the. 
compounds  CH4,  CC14,  CO2,  SiH4,  SiCl4,  SiO2,  etc.  ;  but 
they  may  also  appear  as  bivalent  elements,  as  seen  in  the 
compound  CO.  So  too,  while  the  elements  of  Family  V 
are  quinquivalent,  as  in  PC15,  NH4C1,  etc.,  they  may  also 
be  trivalent,  as  in  PC13,  NH3,  etc. 

What  we  call  valence  does  not  then  appear  to  be  an 


156  INORGANIC  CHEMISTRY. 

unchangeable  property  of  the  elements,  but  a  property 
Avhich  may  change  according  to  conditions.  It  appears 
further  that  a  given  element  may  have  one  valence  tow- 
ards one  element  and  another  valence  towards  another 
element.  This  is  most  strikingly  seen  on  comparing 
the  formulas  of  the  hydrogen  compounds  of  the  ele- 
ments of  Families  V,  VI,  and  VII  with  those  of  their 
oxygen  compounds.  The  members  of  Family  VII  com- 
bine with  hydrogen  in  only  one  proportion,  and  that  is 
the  simplest  possible.  Towards  hydrogen  these  elements 
are  univalent,  and  their  valence  towards  hydrogen  is  con- 
stant. On  the  other  hand,  they  combine  with  oxygen 
and  with  hydroxyl  in  several  proportions,  and  judging  by 
the  composition  of  these  compounds,  the  valence  tow- 
ards oxygen  varies  from  1  to  7.  The  members  of 
Family  VI  are  bivalent  towards  hydrogen,  and  their  hy- 
drogen valence  is  constant ;  but  they  combine  with  oxy- 
gen and  hydroxyl  in  several  proportions,  and  the  compo- 
sition of  the  compounds  indicates  that  their  valence 
towards  oxygen  varies  from  2  to  6.  The  hydrogen  val- 
ence of  the  members  of  Family  V  is  3,  while  the  oxygen 
valence  varies  from  1  to  5.  Finally,  the  hydrogen  val- 
ence of  the  members  of  Family  IV  is  4,  while  the  oxygen 
valence  varies  from  2  to  4.  As  regards  the  hydrogen 
valence  of  the  members  of  Families  I  to  III,  but  little  is 
known.  These  elements  do  not  generally  combine  with 
hydrogen,  though  some  of  them  do.  'Towards  oxygen 
their  valence  is  fairly  constant,  though  some  variations 
are  observed  as  in  the  case  of  copper  and  mercury. 

Judging  then  by  the  composition  of  the  compounds, 
we  are  justified  in  making  a  distinction  between  the 
hydrogen-valence  of  some  elements  and  their  oxygen- 
valence.  While  the  former  is  constant,  the  latter  is  sub- 
ject to  variations.  In  those  cases  in  which  there  is  a 
marked  difference  between  the  hydrogen-valence  of  an  ele- 
ment and  its  maximum  oxygen-valence,  the  maximum 
valence  towards  chlorine  is  greater  than  the  hydrogen- 
valence  and  less  than  the  maximum  oxygen-valence. 
This  is  shown  in  the  case  of  sulphur  ;  the  formulas  of  its 


THE  ELEMENTS  IN  THE  NATURAL  SYSTEM.        157 

hydrogen  compound  and  of  its  highest  compounds  with 
•chlorine  and  oxygen  being  respectively 

SH2,        SC14,        SO,. 

Prom  this  it  appears  that  the  maximum  valence  of  sul- 
phur towards  hydrogen  is  2,  towards  chlorine  4,  and 
towards  oxygen  6. 

Connection  between  the  Position  of  the  Elements  in  the 
Natural  System  and  their  Chemical  Properties. — The 
changes  in  composition  of  the  oxygen  and  hydrogen 
compounds  and  of  the  hydroxides  from  Family  I  to  VII 
have  been  referred  to.  Another  fact  of  great  impor- 
tance is  that  the  elements  of  Group  I  are  the  most 
strongly  marked  base-forming  elements,  while  those  of 
Group  YII  are  the  most  strongly  marked  acid-forming 
elements.  Passing  in  either  direction  the  character  of 
the  elements  becomes  less  pronounced,  until  in  the  mid- 
dle (Group  IV),  elements  which  form  neither  strongly 
marked  acids  nor  strongly  marked  bases  are  found. 
Thus,  beginning  with  sodium,  this  element  forms  a 
strong  base,  magnesium  forms  a  weaker  base,  the  hy- 
droxide of  aluminium  is  a  still  weaker  base.  Beginning, 
on  the  other  hand,  with  chlorine  at  the  other  end  of  the 
same  series,  its  hydrogen  compound  is  a  strongly  marked 
acid  ;  that  of  sulphur  is  an  acid,  but  less  marked  in  char- 
acter than  hydrochloric  acid ;  that  of  phosphorus  has  no 
acid  properties,  nor  has  that  of  silicon.  The  hydroxides 
of  these  four  elements  have  acid  properties.  Each  one, 
however,  forms  several  acids,  and  it  is  difficult  to  com- 
pare them,  as  some  of  those  of  chlorine  are  strongly 
marked  and  others  not,  as  we  have  seen. 

Some  very  interesting  variations  in  properties  are  also 
noticed  in  passing  from  one  end  of  a  group  of  a  natural 
family  to  the  other.  Thus  in  Group  B,  Family  VII,  the 
activity  of  the  elements  grows  less  from  fluorine  to  iodine, 
or,  as  we  commonly  say,  fluorine  is  the  strongest  ele- 
ment in  the  group,  and  then  follow,  in  order,  chlorine, 
bromine,  and  iodine. 


158  INORGANIC  CHEMISTRY. 

The  remarkable  relations  above  referred  to  are  summed 
up  in  the  periodic  law  : 

The  properties  of  an  element  are  periodic  functions  of  the 
atomic  weight. 

It  appears  that  if  an  element  has  a  certain  atomic 
weight  it  must  have  certain  properties,  and  that  if  the 
atomic  weight  is  known  the  properties  can  be  stated, 
just  as,  if  the  properties  are  known,  the  atomic  weight 
can  be  approximately  stated.  When  the  law  was  first 
stated,  Mendeleeff  predicted  the  discovery  of  certain  ele- 
ments to  fill  some  of  the  vacant  places  in  the  table.  At 
that  time  the  elements  gallium,  Ga,  scandium,  Sc,  and 
germanium,  Ge,  were  not  known.  Not  only  was  their 
discovery  predicted,  but  their  properties  were  clearly 
stated  years  before  they  were  brought  to  light.  Within 
the  last  few  years  these  three  elements  have  been  dis- 
covered, and  a  remarkable  agreement  is  observed  be- 
tween their  properties  as  determined  by  observation  and 
as  foretold  by  Mendeleeff  by  the  aid  of  the  periodic 
law. 

The  relations  between  the  atomic  weights  and  proper- 
ties will  appear  more  and  more  clearly  as  our  study  of 
the  elements  proceeds.  The  natural  arrangement  of  the 
elements  suggested  by  the  periodic  law  is  adopted  in 
this  book.  The  elements  hydrogen,  oxygen,  and  chlo- 
rine were  studied  at  the  outset  in  order  to  illustrate  the 
methods  of  studying  chemical  problems,  and  as  exam- 
ples of  chemical  elements  in  general.  It  is,  however, 
now  time  to  take  up  the  elements  systematically,  and  to 
learn  what  may  be  necessary  in  regard  to  them  in  order 
to  get  as  clear  a  notion  as  possible  of  the  facts  and  prin- 
ciples of  the  science  of  chemistry. 

Plan  to  be  followed. — The  most  systematic  method  of 
procedure  in  studying  the  elements  would  be  to  begin 
with  Family  I,  Group  A  (see  Lothar  Meyer's  Table, 
p.  151),  then  to  take  up  Group  B  of  the  same  family  ; 
and  so  on  in  order,  ending  with  Family  VIII.  It  seems 
better,  however,  to  begin  with  Family  VII ;  to  follow 
with  Families  VI,  V,  and  IV ;  and  then  to  take  up  in 
order  Families  I,  II,  III,  and  VIII.  The  main  reason 


THE  ELEMENTS  IN  THE  NATURAL  SYSTEM.        159 

for  this  is  that  it  is  impossible  to  study  most  of  the  mem- 
bers of  Families  I,  II,  III,  and  VIII  without  a  knowl- 
edge of  several  of  the  elements  of  Families  VII,  VI,  V, 
and  IV,  while  these  last  families  can  be  studied  with 
only  slight  reference  to  the  others.  It  is  proposed  then 
to  begin  with  Group  B,  Family  VII,  the  members  of 
which  are  very  much  like  chlorine.  The  only  member 
of  Group  A  of  this  family  is  manganese.  While  man- 
ganese resembles  the  members  of  the  chlorine  group  in 
some  respects,  it  has  other  properties  which  ally  it  to 
the  so-called  base-forming  elements.  So  also  the  mem- 
bers of  Group  A,  Family  VI,  are  like  the  members  of  the 
oxygen  or  sulphur  group,  but  they  are  also  allied  to  the 
base-forming  elements.  A  similar  difference  is  observed 
between  the  members  of  Groups  A  and  B,  Family  V. 

While  the  plan  above  sketched  takes  into  considera- 
tion the  greater  number  of  the  analogies  of  the  elements, 
there  are  other  analogies  which  are  not  brought  out. 
Thus,  as  will  be  seen  in  due  time,  the  elements  alumin- 
ium, chromium,  manganese,  and  iron  are  analogous  in 
some  respects,  but  by  following  the  plan  sketched  they 
will  be  taken  up  in  different  groups.  This  appears  to  be 
justified,  however,  when  we  consider  the  entire  conduct 
of  these  elements,  and  do  not  confine  ourselves  to  a  study 
of  only  a  few  reactions  which,  being  useful  for  some 
purposes,  have  been  studied  more  carefully  than  others 
which  from  a  scientific  point  of  view  are  perhaps  just  as 
important. 


CHAPTER  XII. 

THE    ELEMENTS     OF    FAMILY    VII,    GROUP   B: 
FLUORINE-CHLORINE—BROMINE—IODINE. 

General. — The  elements  of  this  group  are  commonly 
called  the  halogens.  The  best  known  member  of  the 
group  is  chlorine,  which  has  already  been  treated.  Al- 
though fluorine  is  in  general  like  the  other  members  of 
the  group,  it  differs  from  them  in  some  respects,  and  it 
certainly  is  not  as  much  like  them  as  they  are  like  one 
another.  While  chlorine,  bromine,  and  iodine  accom- 
pany one  another  in  nature,  fluorine  compounds  are  not 
generally  found  in  company  with  compounds  of  the  other 
elements  of  the  family.  In  those  cases  in  which  chlorine, 
bromine,  and  iodine  are  found  together,  chlorine  is  gen- 
erally present  in  largest  quantity,  and  iodine  in  smallest 
quantity.  Fluorine  and  chlorine  are  gases  under  ordi- 
nary conditions,  while  bromine  is  a  liquid  and  iodine  is 
a  solid.  Fluorine,  bromine,  and  iodine  form  with  hy- 
drogen the  compounds  hydrofluoric  acid,  HF,  hydro- 
bromic  acid,  HBr,  and  hydriodic  acid,  HI,  which 
are  analogous  to  hydrochloric  acid.  All  these  com- 
pounds are  gases  which  have  marked  acid  properties. 
With  oxygen,  fluorine  does  not  combine,  whereas  chlo- 
rine, bromine,  and  iodine  combine  with  it  in  a  number  of 
proportions,  as  has  already  been  seen  in  the  case  of  chlo- 
rine. Among  themselves  these  elements  also  form  some 
compounds :  thus  bromine  and  chlorine  form  the  com- 
pound BrCl;  iodine  forms  the  compounds  IC1,  IC13, 
IBr,  and  IF5.  It  appears  from  this  that  the  valence  of 
iodine  towards  bromine  is  1,  towards  chlorine  3,  and 
towards  fluorine  5. 

Towards  base-forming  members  the  elements  of  this 
group  are  univalent,  as  shown  in  such  compounds  as 
NaCl,  KBr,  CaCla,  KI,  etc.  They,  however,  appear  to 

(160) 


BROMINE:— OCCURRENCE-PREPARATION.  161 

have  a  valence  greater  than  1  in  some  compounds  known 
as  double  salts.  These  can  be  explained  satisfactorily 
only  by  assuming  that  in  them  the  element  is  in  combi- 
nation with  itself  and  has  a  valence  greater  than  1. 

BROMINE,  Br  (At.  Wt.  79.34). 

Occurrence. — This  element  occurs  in  nature  in  com- 
pany with  chlorine.  Chlorine,  as  has  been  stated,  occurs 
mostly  in  combination  with  sodium,  as  sodium  chloride, 
or  common  salt.  In  several  of  the  great  salt-beds 
bromine  occurs  in  the  form  of  sodium  bromide,  NaBr,  and 
in  some  places  it  occurs  as  magnesium  bromide,  MgBr2. 
The  chief  source  of  bromine  is  the  mother-liquors  from 
the  salt  works.  When  a  solution  containing  a  large 
quantity  of  sodium  chloride  and  a  small  quantity  of  bro- 
mide is  evaporated,  the  chloride  is  first  deposited,  and 
from  the  mother-liquors  the  bromide  mixed  with  chlo- 
ride is  deposited.  The  great  beds  at  Stassfurt  are  par- 
ticularly rich  in  bromides,  and  a  great  deal  of  bromine 
is  made  from  the  salts  which  occur  in  this  locality. 

Preparation. — Bromine  can  be  prepared  from  the 
bromides  in  the  same  way  that  chlorine  is  made  from  the 
chlorides  :  by  first  treating  with  sulphuric  acid,  thus  lib- 
erating hydrobromic  acid,  and  then  treating  with  man- 
ganese dioxide,  or,  better,  by  mixing  the  bromide  with 
manganese  dioxide  and  treating  the  mixture  with  sul- 
phuric acid.  The  reaction  is  represented  by  the  equation 

2NaBr  +  MnO2  +  2H2SO4  = 

Na2SO4  +  MnSO4  +  2H2O  +  Br2. 

Or  it  may  be  represented  as  taking  place  in  different 
stages.  First  the  sulphuric  acid  would  liberate  hydro- 
bromic acid  from  the  bromide,  and  this  would  act  upon 
the  manganese  dioxide  thus  : 

MnO2  +  4HBr  -=  MnBr2  -f  2H20  +  Br2. 

But  sulphuric  acid  would  act  upon  manganous  bromide, 
MnBr2,  thus : 

MnBr3  -f  H2S04  =  MnS04  +  2HBr ; 


162  INORGANIC  CHEMISTRY. 

and  the  hydrobromic  acid  would  then  again  react  with 
manganese  dioxide,  etc. 

Another  method  for  the  preparation  of  bromine  de- 
pends upon  the  fact  that  chlorine  has  the  power  to  set 
bromine  free  from  its  compounds.  If,  therefore,  a  solu- 
tion containing  a  bromide  is  treated  with  manganese 
dioxide  and  hydrochloric  acid,  the  chlorine  which  is 
formed  from  the  hydrochloric  acid  will  act  upon  the 
bromide  and  bromine  will  be  given  off.  This  method  is 
used  at  Stassfurt. 

Properties. — Bromine  is  a  heavy,  dark-red  liquid  at 
ordinary  temperatures.  If  exposed  to  the  air  it  is  con- 
verted into  a  vapor  of  a  brownish-red  color.  It  boils  at 
58-58.6°,  and  at  —  7.3°  it  is  solid.  It  has  an  extremely 
disagreeable  odor,  to  which  fact  it  owes  its  name  (from 
flpcdfiios,  a  stench).  From  carbon  disulphide  at  —  90°  it 
crystallizes  in  fine  dark-red  needles. 

Its  properties  are  similar  to  those  of  chlorine.  It  acts 
violently  upon  organic  substances ;  attacking  the  skin, 
and  the  membranes  lining  the  passages  of  the  throat  and 
lungs.  Wounds  caused  by  the  liquid  coming  in  contact 
with  the  skin  are  painful  and  serious,  and  it  must  there- 
fore be  handled  with  great  care. 

Like  chlorine,  bromine  is  dissolved  by  water,  one  part 
dissolving  in  33.3  parts  at  15°.  The  solution,  which  has 
a  reddish  color  and  the  odor  of  bromine,  is  called  bro- 
mine tvater.  At  a  low  temperature  bromine  forms  with 
water  a  compound  in  every  way  analogous  to  chlorine 
hydrate,  wz.,  bromine  hydrate,  Br2  -|-  10H2O.  This  de- 
composes when  left  in  contact  with  the  air  at  ordinary 
temperatures. 

Chemical  Conduct  of  Bromine. — Bromine  acts  chemi- 
cally like  chlorine.  It  was  pointed  out  that  chlorine 
acts  in  three  different  ways  :  (1)  By  direct  addition  ;  (2) 
by  substitution;  and  (3)  by  liberating  oxygen  from 
water,  as  in  bleaching  and  other  oxidizing  processes. 
Bromine  is  capable  of  acting  in  all  three  ways.  It  com- 
bines directly  with  base-forming  elements  or  metals,  as 
iron,  aluminium,  potassium,  etc.  ;  also  with  the  acid- 
forming  elements,  as  sulphur,  phosphorus,  etc.  It  com- 


HYDROBROMIC  ACID.  163 

bines  with  hydrogen  almost  as  readily  as  chlorine  does. 
With  oxygen  it  does  not  combine  directly,  and  in  this 
respect  also  it  is  like  chlorine. 

It  acts  upon  compounds  containing  hydrogen  almost 
as  readily  as  chlorine  does,  replacing  the  hydrogen  and 
forming  bromine  substitution-products.  Thus  benzene, 
C6HB,  yields  the  products  C6H5Br,  C6H4Br2,  C6H3Br3, 
C6H2Br4,  etc.,  and  the  hydrogen  which  leaves  the  com- 
pound passes  off  in  combination  with  bromine  in  the 
form  of  hydrobromic  acid. 

It  bleaches  like  chlorine,  partly  by  direct  action  and 
disintegration  of  the  organic  dye-stuffs,  partly  by  action 
upon  water,  liberating  oxygen. 

A  solution  of  bromine  in  water  left  exposed  to  the 
direct  sunlight  loses  its  color  and  becomes  acid  in  conse- 
quence of  the  decomposition  of  the  water,  as  in  the  case 
of  chlorine  : 


Br2    +H,O    = 
or       2Bra  +  2H2O  =  4HBr  +  O2. 

Uses  of  Bromine.  —  Bromine  and  its  compounds  are 
used  in  photography,  medicine,  and  to  some  extent  in 
the  manufacture  of  coal-tar  colors.  It  is  manufactured 
in  large  quantity,  and  a  good  proportion  of  it  is  manu- 
factured in  the  United  States.  According  to  the  official 
report  the  production  of  bromine  in  the  United  States 
in  the  year  1896  amounted  to  over  500,000  pounds. 

Hydrobromic  Acid,  HBr.  —  The  only  compound  which 
bromine  forms  with  hydrogen  alone  is  hydrobromic  acid. 
This  is  in  all  respects  very  much  like  hydrochloric  acid. 
It  is  set  free  from  bromides  by  the  action  of  sulphuric 
acid,  but  owing  to  its  instability  it  acts  upon  the  sul- 
phuric acid,  causing  decomposition.  The  elements  hy- 
drogen and  bromine  are  not  held  together  as  firmly  in 
hydrobromic  acid  as  hydrogen  and  chlorine  are  in  hy- 
drochloric acid.  Consequently,  if  hydrobromic  acid  is 
brought  together  with  certain  substances  which  contain 
oxygen  it  gives  up  its  hydrogen  to  the  oxygen.  This  is 


164  INORGANIC  CHEMISTRY. 

seen  in  the  conduct  towards  manganese  dioxide.  But 
towards  this  substance  both  hydrochloric  and  hydro- 
bromic  acids  act  in  essentially  the  same  way.  Sulphuric 
acid  does  not,  however,  give  up  its  oxygen  as  readily  as 
manganese  dioxide,  and  the  difference  in  the  stability  of 
the  hydrogen  compounds  of  chlorine  and  bromine  is 
seen  very  clearly  in  their  conduct  towards  sulphuric 
acid.  Hydrochloric  acid  does  not  act  upon  sulphuric 
acid  at  all.  Hydrobromic  acid  acts  according  to  the 
following  equation : 

2HBr  +  H2SO4  =  2H2O  +  SO2  +  Br2. 

The  action  consists  in  the  decomposition  of  the  hydro- 
bromic  acid  into  bromine  and  hydrogen,  and  the  subse- 
quent action  of  the  nascent  hydrogen  upon  the  sulphuric 
acid  thus : 

2HBr  =  2H  +  Br2 ;  and 
H2S04  +  2H  =  2H20  +  S02. 

The  hydrobromic  acid  acts  here,  then,  as  a  reducing  agent, 
and  the  sulphuric  acid  as  an  oxidizing  agent.  It  is  plain 
that  hydrobromic  acid  cannot  be  made  in  pure  condition 
by  the  action  of  sulphuric  acid  upon  a  bromide.  Some 
of  the  hydrobromic  acid,  to  be  sure,  escapes  the  action 
of  the  sulphuric  acid,  but  at  best  it  is  always  mixed  with 
the  compound  SO2,  or  sulphur  dioxide,  which  is  a  gas, 
and  with  bromine. 

It  can  be  made  by  passing  a  mixture  of  hydrogen  and 
bromine  over  heated  finely  divided  platinum.  An  ap- 
paratus has  been  devised  for  making  hydrobromic  acid 
in  this  way  in  quantity. 

It  can  also  be  made  by  allowing  bromine  to  act  upon 
an  organic  compound  containing  hydrogen.  Substitut- 
ing action  takes  place  and  hydrobromic  acid  is  given  off. 
Thus,  if  a  compound  of  the  formula  C10H22  were  used,  the 
reaction  would  be  represented  in  this  way  : 


HTDROBROMIC  ACID.  165 

The  product  C10H21Br,  or  the  bromine  substitution- 
product,  would  not  be  volatile  at  ordinary  temperatures, 
and  therefore  only  the  hydrobromic  would  be  given  off. 

The  method  most  commonly  adopted  in  the  labora- 
tory consists  in  treating  phosphorus  with  bromine  and 
water.  In  all  probability  the  bromine  acts  first  upon  the 
phosphorus,  forming  the  product  PBr3  or  PBr5  accord- 
ing to  the  proportions  of  the  substances  used.  Both 
these  substances  are  decomposed  by  water,  the  first 
forming  phosphorous  acid  and  hydrobromic  acid,  accord- 
ing to  this  equation : 

(  Br      HHO 

P^  Br  +  HHO  =  PO3H3  +  3HBr, 
( Br      HHO 

or      PBr3     +  3H,0  =  PO3H3  +  3HBr  ; 

the  second  forming  phosphoric  acid  and  hydrobromic 
acid  : 

PBr5  +  4H2O  =  P04H3  +  5HBr. 

The  gas  thus  formed  can  be  freed  from  bromine  by 
passing  it  through  a  tube  containing  phosphorus. 

Properties. — Hydrobromic  acid  is  a  colorless  gas  which 
forms  fumes  in  contact  with  the  air  in  consequence  of  its 
attraction  for  moisture.  It  dissolves  in  water  in  large 
proportion.  The  solution  conducts  itself  much  like  hy- 
drochloric acid.  When  boiled  a  compound  of  definite 
composition  passes  over  under  ordinary  conditions.  It 
corresponds  to  the  formula  HBr  -f-  5H2O,  but  here,  as 
with  the  hydrate  of  hydrochloric  acid,  the  composition 
changes  with  the  pressure.  With  metallic  hydroxides 
or  bases,  hydrobromic  acid  forms  bromides,  as  hydro- 
chloric acid  forms  chlorides : 

KOH  +  HBr  =  KBr  +  H2O. 

Compounds  of  Bromine  with  Hydrogen  and  Oxygen. — 
With  hydrogen  and  oxygen  bromine  forms  compounds 
which  closely  resemble  those  which  chlorine  forms  with 
the  same  elements.  They  are :  Hypobromous  acid, 


166  INORGANIC  CHEMISTRY. 

HBrO ;  bromic  acid,  HBrO3 ;  and  perhaps  perbromic 
acid,  HBrO4. 

Hypobromom  acid,  HBrO,  is  made  by  reactions  which 
are  entirely  analogous  to  those  used  in  making  hypo- 
chlorous  acid.  When  bromine  acts  upon  a  dilute  solu- 
tion of  sodium  or  potassium  hydroxide,  reaction  takes 
place  thus : 

2KOH  +  Bra  =  KBr  +  KBrO  +  H9O. 

So  also  bromine  vapor  acting  upon  slaked  lime  or  cal- 
cium hydroxide  forms  a  compound  similar  to  bleaching 
powder.  Hypobromous  acid  has  not  been  prepared  in 
pure  condition  owing  to  its  instability. 

Bromic  acid,  HBrO3,  is  not  known  in  pure  condition. 
Its  salts  are  made  in  the  same  way  as  the  chlorates 
are ;  the  principal  reaction  made  use  of  for  the  purpose 
being  that  between  bromine  and  concentrated  potassium 
hydroxide : 

3Bra  +  6KOH  =  5KBr  +  KBrO3  +  3H2O. 

The  decompositions  of  the  bromates  are  much  like 
those  of  the  chlorates. 

As  regards  the  existence  of  perbromic  acid  there  is 
some  doubt.  It  is  stated  by  one  observer  that  he  ob- 
tained it  by  treating  perchloric  acid  with  bromine : 

HC1O4  +  Br  =  HBr04  +  01. 

Others  have  not  succeeded  in  getting  it  in  this  or  in  any 
other  way. 

Compound  of  Bromine  and  Chlorine. — When  chlorine 
is  passed  into  liquid  bromine  it  is  absorbed  in  large 
quantity.  If  the  process  is  carried  on  at  a  low  tempera- 
ture the  product  BrCl  is  formed.  Above  10°  it  under- 
goes decomposition.  Although  it  is  unstable,  there  is 
no  good  reason  for  regarding  this  substance  as  anything 
but  a  chemical  compound.  There  are  many  chemical 
compounds  known  which  are  less  stable  than  this. 


IODINE:—  OCCURRENCE-PREPARATION".  167 

IODINE,  I  (At.  Wt.  125.89). 

Occurrence. — Iodine,  as  has  already  been  stated,  occurs 
in  company  with  chlorine  and  bromine  in  nature,  but  in 
smaller  quantity  than  these.  The  relative  quantity  in 
sea  water  is  extremely  small.  The  sea  plants,  however, 
assimilate  it,  and  the  ashes  of  these  plants  contain  a 
considerable  quantity  of  compounds  of  iodine.  It  also 
occurs  in  small  quantity  in  the  great  beds  of  soda  salt- 
peter, or  sodium  nitrate,  which  are  found  in  Chili,  South 
America.  It  occurs  in  small  quantity  in  combination 
with  silver,  and  also  in  combination  with  lead  and  with 
mercury. 

Preparation. — The  method  of  obtaining  iodine  from 
its  salts  is  like  that  used  in  making  chlorine  and  bromine 
from  the  chlorides  and  bromides.  It  consists  in  treat- 
ing the  iodides  with  sulphuric  acid  and  manganese 
dioxide. 

2KI  +  MnO8  +  2H2SO4  =  K2SO4  +  MnSO4  +  2H20  + I2. 

The  iodine,  although  solid  at  the  ordinary  tempera- 
ture, is  easily  volatilized,  and  if  the  mixture  mentioned 
is  heated,  iodine  vapor  passes  over  and  may  be  con- 
densed in  appropriately  arranged  vessels. 

On  the  large  scale  iodine  is  obtained  mostly  from  sea- 
weed. On  the  coasts  of  Scotland,  Ireland,  and  France 
the  sea- weed  which  is  thrown  up  by  storms  is  gathered, 
dried,  and  burned.  The  organic  portions  are  thus  de- 
stroyed, and  the  mineral  or  earthy  portions  are  left  be- 
hind as  ashes.  This  incombustible  residue  is  called 
kelp.  It  contains  a  small  percentage  of  potassium 
iodide,  from  .5  to  2  per  cent  according  to  the  sea- weed 
used.  The  dried  weed  was  formerly  burned  in  cavities 
dug  in  the  earth,  but  of  late  years  the  process  has  in 
some  places  been  much  improved,  and  the  yield  in  kelp 
increased. 

In  Scotland  the  iodine  is  liberated  by  means  of  sul- 
phuric acid  and  manganese  dioxide.  In  France,  how- 


168  INORGANIC  CHEMISTRY. 

ever,  this  is  effected  by  passing  chlorine  into  the  solution 
containing  the  iodide.  If  too  little  chlorine  is  used  all 
the  iodine  is  not  separated ;  if  too  much,  a  compound  of 
iodine  and  chlorine  is  formed,  or  an  iodate,  in  conse- 
quence of  the  oxidizing  action  of  the  chlorine  on  the 
iodine. 

The  iodine  which  occurs  in  Chili  saltpeter,  NaNO3,  is 
in  the  form  of  sodium  iodate,  NaIO3,  and  iodide,  Nal, 
and  to  some  extent  as  magnesium  iodide,  MgI2.  Most 
of  the  iodine  now  in  the  market  is  made  from  this  ma- 
terial, and  the  competition  created  in  this  way  has  led 
to  a  careful  study  of  the  process  for  obtaining  iodine 
from  kelp.  Sea- weed  is  now  collected  from  certain  parts 
of  the  ocean  where  it  grows  in  large  quantity,  vessels 
being  sent  out  for  the  purpose. 

Properties. — Iodine  is  a  grayish-black  crystallized 
solid.  At  ordinary  temperatures  it  is  volatile.  Accord- 
ing to  the  most  reliable  determinations  it  melts  at 
113-115°,  and  boils  at  250°.  The  vapor  has  a  violet 
color  when  mixed  with  air.  When  in  pure  condition 
it  is  intensely  blue.  At  temperatures  considerably 
above  the  boiling  point  the  specific  gravity  of  iodine 
vapor  is  such  as  to  show  that  its  molecular  weight  is 
approximately  254,  or  twice  the  atomic  weight.  As  the 
temperature  is  raised,  however,  the  specific  gravity  is 
lowered,  until,  finally,  at  a  very  high  temperature,  it 
becomes  about  half  what  it  is  at  lower  temperatures. 
This  is  accounted  for  by  supposing  that  at  the  lower 
temperatures  the  molecules  of  iodine  consist  of  two 
atoms  each,  while  as  the  temperature  is  raised  these 
molecules  are  gradually  broken  down,  so  that  at  the 
temperature  at  which  the  lowest  specific  gravity  is 
reached  the  iodine  vapor  consists  of  free  atoms,  or  the 
atoms  and  molecules  are  then  identical,  and  the  specific 
gravity  is  therefore  only  half  what  it  is  when  the  mole- 
cules consist  of  two  atoms. 

Iodine  has  a  characteristic  strong  taste.  It  acts  upon 
the  mucous  membranes,  but  much  less  energetically 
than  chlorine  or  bromine.  It  colors  the  skin  yellowish- 


HYDRIODIC  ACID.  169 

brown,  and  acts  as  an  absorbent,  causing  the  reduction 
of  some  kinds  of  swellings. 

It  dissolves  slightly  in  water,  easily  in  alcohol,  and 
easily  in  a  water  solution  of  potassium  iodide.  The 
solution  in  alcohol  is  known  as  tincture  of  iodine.  It 
-dissolves  also  in  carbon  disulphide,  CS2,  and  in  chloro- 
form forming  solutions  which  have  a  beautiful  deep 
violet  color. 

In  general,  iodine  conducts  itself  chemically  like  bro- 
mine and  chlorine,  only  it  acts  in  almost  all  reactions 
less  energetically  than  the  other  two  elements.  It  com- 
bines directly  with  a  number  of  elements,  as  with  hydro- 
gen, sulphur,  phosphorus,  iron,  mercury,  etc.  In  pres- 
ence of  water  it  acts  as  an  oxidizer  just  as  chlorine  and 
bromine  do,  but  less  energetically.  Thus  it  oxidizes 
sulphurous  acid,  H2SO3,  to  sulphuric  acid,  H2SO4 : 

H2S03  +  I2  +  H2O  =  H2S04  +  2HI. 

As  a  substituting  agent  it  does  not  act  as  readily  as 
chlorine  and  bromine,  though  iodine  substitution-prod- 
ucts are  made  in  large  quantities,  particularly  in  con- 
nection with  the  manufacture  of  dye-stuffs. 

Iodine  is  used  extensively  in  the  dye-stuff  industry,  in 
photography,  and  in  medicine.  One  factory  in  Scotland 
makes  on  an  average  60  tons  of  iodine  a  year. 

Hydriodic  Acid,  HI. — The  affinity  of  hydrogen  for 
iodine  is  less  than  for  bromine,  and  therefore  hydriodic 
acid  cannot  be  made  pure  by  treating  an  iodide  with 
sulphuric  acid.  The  hydrogen  of  the  hydriodic  acid 
acts  upon  the  sulphuric  acid  very  readily,  and  according 
to  the  conditions  the  following  reactions  may  take 
place : 

H2S04  +  SHI  =  4H20  +  SH3  +  41, ; 
HaS04  +  6HI  =  4H20  +  S  +  31, ; 
H2S04  +  2HI  =  2H20  +  S02  + I2. 

On  treating  potassium  iodide  with  sulphuric  acid, 
therefore,  there  may  be  formed,  in  addition  to  hydriodic 


170  INORGANIC  CHEMISTRY. 

acid  and  free  iodine,  sulphur  dioxide,  sulphur  and  hydro- 
gen sulphide. 

The  method  adopted  for  the  preparation  of  hydriodic 
acid  is  like  that  used  for  the  preparation  of  hydrobromic 
acid.  It  consists  in  treating  phosphorus  with  iodine  and 
water.  The  reactions  involved  are  of  the  same  kind  as 
those  which  were  discussed  under  hydrobromic  acid. 
The  iodine  probably  acts  at  first  on  the  phosphorus, 
forming  a  compound  of  phosphorus  and  iodine,  which 
then  in  turn  is  decomposed  by  the  water.  The  reactions 
which  generally  take  place  are  those  represented  by  the 
following  equations : 

P     +31       =  PI8; 

PI3  +  3HaO  =  P03H3  +  SHI. 

Hydriodic  acid  is  a  colorless  transparent  gas  like  hy- 
drochloric and  hydrobromic  acids.  It  also  like  these  dis- 
solves in  water  in  large  quantity,  and  when  brought  in 
contact  with  the  air  it  forms  dense  white  fumes.  When 
boiled  the  water  solution  conducts  itself  like  those  of 
hydrochloric  and  hydrobromic  acids.  The  liquid,  which 
boils  at  127°  under  the  ordinary  atmospheric  pressure, 
contains  57  per  cent  hydriodic  acid.  If  the  solution  of 
the  gas  in  water  is  allowed  to  stand,  decomposition  be- 
gins in  consequence  of  the  action  of  the  oxygen  of  the 
air.  The  hydrogen  is  oxidized  to  water  and  the  iodine 
is  set  free,  coloring  the  solution  brown. 

When  heated,  the  gas  begins  to  decompose  at  180°, 
and  at  higher  temperatures  the  decomposition  takes 
place  rapidly.  The  products  are  simply  hydrogen  and 
iodine.  In  consequence  of  the  ease  with  which  hydrio- 
dic acid  breaks  down,  yielding  free  hydrogen,  it  is  an  ex- 
cellent reducing  agent,  and  it  is  frequently  used  in  the 
laboratory  for  the  purpose  of  extracting  oxygen  from 
substances.  Its  action  upon  sulphuric  acid  has  already 
been  spoken  of.  The  reason  why  it  acts  so  well  is  that 
the  hydrogen  is  separated  from  the  iodine  with  little  ex- 
penditure of  energy,  and  the  hydrogen  thus  separated 
is  in  the  nascent  state,  or,  as  is  believed,  in  the  atomic 
state. 


IODIC  ACID.  171 

lodic  Acid,  HIO3. — This  compound  is  strictly  analo- 
gous to  chloric  and  bromic  acids,  but  differs  from  them 
in  being  much  more  stable.  It  can  be  made  by  treat- 
ing iodine  with  strong  oxidizing  agents,  as,  for  example, 
concentrated  nitric  acid.  It  is  also  formed  very  easily 
by  passing  chlorine  through  water  in  which  iodine  is 
suspended,  when  hydrochloric  acid  and  iodic  acid  are 
formed,  as  represented  in  this  equation : 

I2  +  5C12  +  6H2O  =  2HI03  +  10HC1. 

The  reaction  is  probably  somewhat  more  complicated 
than  it  appears  from  this  equation,  for  when  chlorine 
acts  upon  iodine  a  compound  of  the  two  elements  is 
first  formed.  Iodine  trichloride  is  decomposed  by  water 
thus  : 

2IC13  +  3H2O  =  5HC1  +  HIO3  +  101. 

Iodine  monochloride  is  also  decomposed  by  water,  giv- 
ing iodic  acid,  hydrochloric  acid,  and  free  iodine  : 

10IC1  +  6H2O  =  10HC1  +  2HIO3  +  4I2. 

Whether  these  chlorides  of  iodine  are  formed  or  not, 
the  prime  causes  of  the  formation  of  iodic  acid  when 
chlorine  acts  upon  iodine  in  water  are  the  oxidizing 
power  of  the  chlorine  and  the  affinity  of  iodine  for 
oxygen. 

When  iodine  is  dissolved  in  an  alkali  the  reaction 
which  takes  place  is  the  same  as  that  which  takes  place 
with  chlorine  and  bromine  under  like  circumstances.  A 
mixture  of  the  iodide  and  iodate  is  formed : 

6KOH  +  3I2  =  5KI  +  KIO3  +  3H2O. 

Iodic  acid  is  a  crystallized  solid,  which  when  heated 
to  170°  loses  water  and  is  converted  into  iodine  pent- 
oxide,  I2O5 : 

2HI03  =  1,0.  +  H2O. 


172  INORGANIC  CHEMISTRY. 

Its  salts  have  the  general  formula  MIO3,  though  it  alsa 
forms  salts  MH(IO3)2  and  MH2(IO3)3.  It  gives  up  its 
oxygen  readily  and  is  therefore  a  good  oxidizing  agent, 
just  as  hydriodic  acid  is  a  good  reducing  agent. 

Iodine  Pentoxide  or  lodic  Anhydride,  I2Oa. — This  com- 
pound is  formed,  as  was  stated  in  the  last  paragraph, 
by  heating  iodic  acid  to  170°.  It  is  a  white  solid  which 
is  easily  soluble  in  water,  forming  iodic  acid.  It  is  de- 
composed when  heated  to  300°.  It  will  be  observed, 
therefore,  that  this  compound  of  iodine  and  oxygen  is 
very  much  more  stable  than  any  of  the  compounds  of 
chlorine  or  bromine  and  oxygen  ;  and  it  is  interesting  to 
note  that  as  the  affinity  for  oxygen  increases,  that  for 
hydrogen  decreases.  In  the  group  chlorine,  bromine, 
and  iodine,  chlorine  has  the  strongest  affinity  for  hydro- 
gen and  the  weakest  for  oxygen,  while  iodine  has  the 
strongest  affinity  for  oxygen  and  the  weakest  for  hydro- 
gen. We  shall  presently  see  that  fluorine,  which  does 
not  unite  with  oxygen,  has  a  stronger  affinity  for  hydro- 
gen than  chlorine  has. 

Anhydrides,  or  Acidic  Oxides. — An  oxide  which,  like 
iodine  pentoxide,  forms  an  acid  when  dissolved  in  water, 
or  which  forms  salts  by  treatment  with  basic  hydroxides, 
is  called  an  anhydride  or  acidic  oxide.  The  oxides  of  the 
base-forming  elements  form  bases  when  dissolved  in 
water,  and  they  are,  therefore,  called  basic  oxides.  As 
examples  of  acidic  oxides  or  anhydrides,  there  may  be 
mentioned  besides  iodic  anhydride,  sulphuric  anhydride, 
SO3 ;  sulphurous  anhydride,  SO2 ;  phosphoric  anhy- 
dride, P2O6 ;  carbonic  anhydride,  CO2.  When  dissolved 
in  water  these  oxides  are  converted  into  acids  as  repre- 
sented in  these  equations  : 

S03  +  H20  =  H2S04 ; 
S02  +H20  =  H2S08; 
P205  +  H20  =  2HP03 ; 
C02  +  H20  =  H2C03. 

Silicic  anhydride,  SiO2,  is  an  example  of  an  acidic  oxide 
which  does  not  dissolve  in  water,  but  which  does  form 
salts  when  treated  with  basic  hydroxides : 


PERIODIC  ACID.  173 

SiO2  +  2KOH  =  K2Si03  +  HaO. 

As  examples  of  basic  oxides  or  oxides  which  when 
treated  with  water  yield  bases,  the  following  may  be 
taken  :  calcium  oxide,  CaO ;  potassium  oxide,  K2O  ;  ba- 
rium oxide,  BaO.  As  has  already  been  shown,  when 
treated  with  water  these  are  respectively  converted  into 
calcium  hydroxide,  Ca(OH)2 ;  potassium  hydroxide, 
KOH;  and  barium  hydroxide,  Ba(OH)2.  There  are, 
however,  many  basic  oxides  which  do  not  dissolve  in 
water,  but  which,  nevertheless,  have  the  power  to  neu- 
tralize acids  and  form  salts.  This  is  true,  for  example, 
of  aluminium  oxide,  A12O3,  lead  oxide,  PbO,  manganous 
oxide,  MnO,  cupric  oxide,  CuO,  etc.  The  action  of  such 
oxides  upon  acids  takes  place  as  represented  below  : 

A1203  +  3H2S04  =  A12(S04)3  +  3H20  ; 
PbO  +  2HN03  =  Pb(N03)2  +  H20  ; 
MnO  +  2HC1     =  MnCl2       +  H2O  ; 
CuO  +  H2S04    =  CuS04      +  H2O, 

Periodic  Acid,  H5IO6. — This  acid  is  analogous  to  per- 
•chloric  acid.  Its  salts  are  formed  by  oxidation  of  iodates 
or  by  heating  iodates,  just  as  perchlorates  are  formed  by 
heating  chlorates.  The  simplest  way  to  make  a  peri- 
odate  is  to  pass  chlorine  into  a  solution  containing  so- 
dium hydroxide  and  sodium  iodate,  when  a  reaction 
takes  place  which  is  at  least  partly  represented  by  the 
following  equation  : 

NaIO3  +  3NaOH  +  C12  =  Na2H3IO.  +  2NaCl. 

The  salt  Na2H3IO6  is  difficultly  soluble  in  water,  and 
therefore  separates  from  the  solution.  From  the  sodium 
salt  the  corresponding  silver  salt,  Ag2H3IO6,  can  be  ob- 
tained, and  when  this  silver  salt  is  treated  with  nitric 
acid  it  is  converted  into  the  simpler  salt,  AgIO4,  which 
is  evidently  derived  from  the  simpler  acid,  HIO4  : 

2A&H.IO.  +  2HNOS  At  2AgNO3  +  4HSO  +  2AgIO.. 


174  INORGANIC  CHEMISTRY. 

The  acid  when  separated  from  its  solutions  is  a  crys- 
tallized solid  which  has  the  composition  H6IO6.  When 
heated  it  undergoes  decomposition,  losing  water  and 
oxygen,  and  yielding  iodic  acid.  It  cannot,  however,  bo 
converted  into  a  compound  of  the  composition  HIO4, 
for  the  loss  of  water  is  always  accompanied  by  a  loss  of 
oxygen.  Like  iodic  acid,  periodic  acid  is  a  good  oxidiz- 
ing agent  in  consequence  of  the  ease  with  which  it  gives 
up  its  oxygen. 

Periodates. — Periodic  acid  yields  a  large  number  of 
salts  the  connection  between  which  and  the  acid  does 
not  appear  clear  at  first  sight.  A  few  examples  will 
suffice  for  the  present  purpose  :  KIO4,  Na5IO6,  Ag3IO5, 
AgJA,  Zn,I,On. 

Constitution  of  Periodic  Acid. — The  complicated  salts 
of  periodic  acid  are  apparently  inexplicable  on  any  other 
theory  than  that  they  are  derived  from  acids  which  are 
closely  related  to  the  hypothetical  acid  I(OH)7.  This  i» 
now  commonly  regarded  as  normal  periodic  acid.  It, 
however,  breaks  down  into  the  ordinary  form  of  the  acid 
by  loss  of  water.  The  relation  is  expressed  thus  : 


H20. 


The  salts  Na2H3IO6,  Na5IO6,  and  others  of  the  same  kind 
are  derived  from  this  acid  by  replacement  of  one  or  more 
of  the  hydrogen  atoms  by  metallic  elements.  The  acid 
of  the  formula  H5IO6  can  also  be  imagined  to  break 
down  into  H3IO5  and  water  thus  : 


,. 

roH 

OH 
OH 
OH  =  I 
OH 
OH 
OH 

OH 
OH 
OH 
OH 
OH 

CONSTITUTION  OF  PERIODIC  ACID. 


175 


The  salt  Ag3IO5  and  similar  known  salts  are  plainly 
derived  from  this  hypothetical  acid  H3IO5.  Finally,  the 
acid  H3IO6  can  also  be  imagined  to  break  down  into 
HIO4  and  water  thus  : 


=  I 


O 


OH 


and  the  salts  like  KIO4  are  derived  from  this  hypo- 
thetical acid.  It  appears,  therefore,  that  the  assump- 
tion of  the  fundamental  normal  acid,  I(OH),,  is  com- 
petent to  explain  the  existence  of  the  salts  which  are 
derived  from  the  acids  H5IO6,  H3IO5,  and  HIO4.  More 
complicated  acids  can  be  formed  by  the  loss  of  water 
from  two  or  more  molecules  of  any  one  of  these  simpler 
acids.  Thus,  if  from  two  molecules  of  the  acid  H3IO6 
one  molecule  of  water  is  taken,  an  acid  of  the  formula 
H4I2O9  would  be  formed  ;  or  if  two  molecules  of  the  acid 
H6IO6  lose  one  molecule  of  water,  the  acid  H8I2On  would 
be  formed.  These  relations  are  made  clear  by  the 
equations  here  given : 


21  < 


O 

O 

OH  =  (HO)202I-0-I02(OH)2  +  H20  ; 

OH 

OH 


'O 
OH 


21  H 


-  (HO)4OI-0-IO(OH)4 


OH 
OH 


A  salt  of  the  formula  Ag4I2O9,  and  another  of  the  for- 
mula Zn4I2On,  are  known.  The  former  is  derived  from 
the  acid  H4I2O8,  the  latter  from  the  acid  H6I,On,  by  the 
substitution  of  four  bivalent  atoms  of  zinc  for  the  eight 
atoms  of  hydrogen.  There  are  many  more  complicated 


176  INORGANIC  CHEMISTRY. 

salts  than  those  mentioned,  but  they  can  all  be  satisfac- 
torily explained  by  the  assumption  that  they  are  related 
to  the  normal  acid  I  (OH)7,  in  which  iodine  is  septivalent. 
The  existence  of  the  periodates,  the  ease  with  which  they 
can  be  explained  by  the  above  method,  and  the  apparent 
impossibility  of  explaining  them  on  the  assumption  that 
iodine  is  univalent,  form  an  exceedingly  strong  argument 
in  favor  of  the  view  that  iodine  is  septivalent  in  these 
compounds. 

Constitution  of  Iodic  Acid  and  the  Oxygen  Acids  of 
Bromine. — The  conclusion  reached  in  regard  to  the  con- 
stitution of  periodic  acid  makes  it  appear  highly  probable 
that  perchloric  acid  has  a  similar  constitution,  and  this 
view  is  now  commonly  accepted,  as  was  stated  when  the 
acid  was  discussed.  Applying  a  similar  method  to  iodic 
acid,  it  appears  probable  that  this  is  derived  from  the  acid 
I(OH)B  by  loss  of  water  : 

§     =    I02(OH)  +  2H20;  or 

rOH 
OH  (O 

!<!  OH    =    I-J  O       +  2HaO. 

OH          (OH 

OH 

The  iodine  is  regarded  as  quinquivalent  in  both  forms  of 
the  acid.  This  is  represented  in  the  case  of  the  acid 
O2I(OH)  by  the  structural  formula 


in  ; 

The  corresponding  compound  of  bromine  is  regarded  as- 
having  the  same  constitution  as  the  iodine  compound. 

Compounds  of  Iodine  with  Chlorine. — When  chlorine 
is  passed  over  dry  crystallized  iodine  it  is  absorbed,  and 
a  compound  of  the  formula  IC1  is  formed.  This  is  a  thick 
reddish-brown,  very  volatile  liquid.  Under  proper  con- 
ditions it  solidifies  in  crystals.  Iodine  chloride  is  decom- 
posed by  water,  the  products  being  iodic  acid,  hydro- 


FL  UORINE:— OCCURRENCE- PROPERTIES.  17? 

chloric  acid,  and  free  iodine,  as  stated  under  lodic  Acid 
(p.  171). 

If  the  passage  of  chlorine  over  iodine  is  continued  be- 
yond the  point  required  for  the  formation  of  the  simple 
compound  IC1,  the  trichloride  IC13  is  formed.  This  is 
a  crystallized  compound  of  a  yellow  color.  When  heated 
it  breaks  down  into  chlorine  and  iodine  monochloride. 
When  treated  with  water  it  is  partly  dissolved  without 
decomposition,  but  it  is  partly  decomposed,  yielding  iodic 
acid,  iodine  monochloride,  and  hydrochloric  acid. 

Compound  of  Iodine  with  Bromine. — There  is  only  one 
compound  of  iodine  and  bromine  known,  and  that  is  the 
one  having  the  formula  IBr.  It  is  a  crystallized  com- 
pound which  is  formed  by  direct  combination  of  the  two 
elements.  It  is  decomposed  by  heat  and  by  water. 

FLUORINE,  F  (At.  Wt.  18.91). 

Occurrence. — This  element  occurs  in  large  quantity  in 
nature,  and  is  widely  distributed,  but  it  is  always  in  com- 
bination with  other  elements.  It  is  found  chiefly  in  com- 
bination with  calcium,  as  fluor-spar  or  calcium  fluoride, 
CaF2,  and  in  combination  with  sodium  and  aluminium,  as 
cryolite,  a  mineral  which  occurs  abundantly  in  Greenland 
and  has  the  composition  represented  by  the  formula 
Na3AlF6  or  AlF8.3NaF.  It  is  called  fluorine  from  the  fact 
that  it  occurs  in  fluor-spar,  which  in  turn  receives  its 
name  for  the  reason  that  it  melts  when  heated  and  is 
therefore  used  as  a  flux  in  heating  chemical  substances 
together  (from  fluo,  I  flow).  On  account  of  the  remark- 
able affinity  of  fluorine  for  other  elements,  all  attempts 
to  prepare  it  in  the  free  condition  failed  until  a  few  years 
ago,  when  its  isolation  was  effected  by  passing  an  electric 
current  through  liquid  hydrofluoric  acid  contained  in  a 
platinum  vessel. 

Properties. — Fluorine  is  a  light  greenish-yellow  gas, 
that  has  recently  been  converted  into  a  liquid  at  a  very  low 
temperature.  It  acts  upon  almost  all  substances.  Thus, 
it  decomposes  water,  yielding  ozone  and  hydrofluoric 
acid  ;  it  combines  directly  with  hydrogen  at  the  ordinary 


178  INORGANIC  CHEMISTRY. 

temperature;  and  with  sulphur,  phosphorus,  iron,  etc., 
with  evolution  of  light  and  heat.  It  does  not,  however, 
act  upon  platinum.  Owing  to  its  active  properties  it  is 
of  course  a  difficult  matter  to  isolate  and  preserve  it. 

Hydrofluoric  Acid,  HF. — Hydrofluoric  acid  is  made 
by  treating  a  fluoride  with  sulphuric  acid.  Thus,  when 
calcium  fluoride  or  fluor-spar  is  used,  this  reaction  takes 
place  : 

CaF2  +  H2S04  =  CaSO4  +  2HF. 

The  reaction  must  be  performed  in  vessels  of  platinum 
or  lead,  as  glass  is  disintegrated  by  the  acid.  In  perfectly 
pure  anhydrous  condition  it  can  be  obtained  by  heating 
the  pure  dry  salt  KHF2,  known  as  acid  potassium  fluor- 
ide. It  is  a  liquid  which  boils  at  19.4°  and  does  not  so- 
lidify even  at  a  very  low  temperature.  The  pure  dry  acid 
in  the  liquid  form  does  not  act  upon  glass.  It  does  not 
clissolve  the  acid-forming  elements,  but  does  dissolve 
most  of  the  base-forming  elements  with  evolution  of  hy- 
drogen and  formation  of  fluorides.  The  gas  acts  upon 
the  skin,  causing  swellings  and  violent  pains.  Inhaled  it 
is  poisonous.  To  preserve  it,  vessels  of  platinum  or  caout- 
chouc must  be  used.  In  the  moist  condition  it  attacks 
glass,  converting  the .  silicon  into  the  fluoride,  SiF4,  and 
the  metals  into  their  fluorides.  A  silicate  of  the  formula 
CaSiO3  would  undergo  the  changes  represented  in  the 
following  equation : 

CaSiO3  +  6HF  =  CaF2  +  SiF4  +  3H2O. 

Silicon  fluoride  is  a  gas,  and  calcium  fluoride  is  soluble 
in  acids.  Thus  calcium  silicate,  which  is  insoluble  in 
water,  is  so  changed  by  hydrofluoric  acid  as  to  be  ren- 
dered soluble.  In  a  similar  way  glass,  which  is  a  com- 
pound resembling  calcium  silicate,  is  rendered  soluble, 
or  is,  as  we  commonly  say,  dissolved,  by  hydrofluoric 
acid. 

When  an  aqueous  solution  of  hydrofluoric  acid  is  boiled 
it  passes  over  at  120°,  and  the  distillate  contains  36  to  38 
per  cent  of  the  acid. 


HYDROFLUORIC  ACID— THE  FLUORIDES.  179 

Hydrofluoric  acid  is  used  for  the  purpose  of  etching 
glass,  particularly  for  marking  scales  on  thermometers 
and  other  graduated  glass  instruments.  The  glass  is 
covered  with  a  thin  layer  of  wax  or  paraffin  and,  at  the 
places  where  the  etching  is  wanted,  marks  are  made 
through  the  paraffin,  so  that  the  glass  is  exposed.  Those 
parts  of  the  glass  which  are  covered  are  not  acted  upon 
by  the  hydrofluoric  acid,  while  those  parts  which  are  not 
covered  are  corroded  and,  when  the  paraffin  is  removed, 
permanent  marks  are  found  corresponding  to  those  made 
through  the  paraffin.  A  solution  of  hydrofluoric  acid  in 
water  is  manufactured  and  sold  in  rubber  bottles. 

The  specific  gravity  of  hydrofluoric  acid  gas  at  about 
100°  leads  to  the  molecular  weight  corresponding  to  the 
formula  HF,  fluorine  having  the  atomic  weight  19.  At 
about  30°  the  specific  gravity  corresponds  to  the  for- 
mula H2F2.  At  lower  temperatures  the  molecular  weight 
appears  to  be  still  greater.  In  solutions  of  ordinary 
concentration  in  water  the  substance  appears  to  have 
the  formula  H2F2,  while  when  the  solution  is  much 
diluted  the  formula  becomes  HF. 

Constitution  of  Hydrofluoric  Acid  and  the  Fluorides.— 
Hydrofluoric  acid  forms  two  series  of  salts  correspond- 
ing to  the  two  general  formulas  MHF2  and  M2F2  or  MF. 
The  former,  of  which  the  salt  KHF2  is  an  example,  are 
called  acid  fluorides,  the  latter  simply  fluorides.  The 
fluorides  are  commonly  represented  by  the  simpler 
general  formula  MF,  though  it  appears  probable  that 
the  doubled  formula  is  more  correct.  It  will  be  seen 
later  that  fluorine  forms  a  large  number  of  so-called 
double  salts  or  double  fluorides,  which  it  is  difficult  to 
explain  in  any  other  way  than  that  they  are  derived  from 
the  acid  H2F2.  Thus  cryolite,  to  which  reference  has 
been  made,  is  called  a  double  fluoride  of  aluminium  and 
sodium,  and  is  generally  expressed  by  the  formula 
A1F3 .  3NaF,  which  means  simply  that  in  some  way  alu- 
minium fluoride  is  combined  with  three  molecules  of 
sodium  fluoride  ;  but  it  is  difficult  to  see  how  this  union 
can  be  effected  without  assuming  that  fluorine  has  a 
greater  valence  than  one.  If  hydrofluoric  acid  has  the 


180  INORGANIC  CHEMISTRY. 

formula  H2F2,  its  constitution  is  probably  this:  H-F-F-H; 
or  possibly  H-F=F-H.  In  the  one  case  the  fluorine 
is  represented  as  bivalent,  in  the  other  as  trivalent,  but 
we  have  no  evidence  in  favor  of  either  view  as  opposed 
to  the  other.  Still  it  is  generally  observed  that  when 
the  valence  of  an  element  varies,  it  changes  from  odd  to 
odd  or  from  even  to  even.  Thus  in  the  case  of  the 
oxygen  acids  of  chlorine,  it  appears  that  the  valence  of 
chlorine  varies  from  1  to  3  to  5  to  7.  Similarly  the 
valence  of  sulphur  varies  from  2  to  4  to  6,  etc.  For  this 
reason  the  view  that  fluorine  is  trivalent  in  hydrofluoric 
acid  is  perhaps  to  be  preferred  to  the  simpler  view  that 
it  is  bivalent.  If  then  the  constitution  of  hydrofluoric 
acid  be  expressed  thus,  H-F=F-H,  the  formation  of  the 
so-called  double  fluorides  is  not  difficult  to  understand. 
The  double  fluoride  above  referred  to,  viz.,  cryolite,  has 
probably  the  constitution  represented  bv  the  formula 

/F-F-Na 
Al^— F=F-Na,  and  the  other  double  fluorides  are  to  be 

\F=F-Na 
regarded  as  having  a  similar  constitution. 

Compound  of  Fluorine  with  Iodine. — The  only  com- 
pound of  fluorine  with  the  members  of  the  chlorine  group 
is  iodine  pentafluoride,  IF6.*  This  is  a  liquid  which  is 
formed  by  treating  silver  fluoride,  AgF,  with  iodine. 
Water  decomposes  it,  forming  iodic  acid  : 

IF5  +  3H2O  =  5HF  +  HIO3. 

Considering  the  compounds  which  the  halogens  form 
with  one  another,  it  appears  that  iodine  combines  with 
bromine  to  form  the  compound  IBr,  with  chlorine  it 
forms  IC13,  and  with  fluorine  IF5 ;  or  its  valence  towards 
bromine  is  1,  towards  chlorine  3,  and  towards  fluorine 
5.  The  farther  removed  in  the  series  the  element  is 
from  iodine  the  greater  is  the  valence  of  iodine  for  it. 

*  There  seems  to  be  some  doubt  in  regard  to  the  existence  of  this 
compound. 


RELATIVE  AFFINITIES  OF  THE  CHLORINE  GROUP.     181 

Tabular  Presentation  of  the  Compounds  of  the  Members 
of  the  Chlorine  Family  with  Hydrogen,  with  Oxygen,  with 
Hydrogen  and  Oxgen,  and  with  One  Another. 

Compounds  with  Hydrogen. 
HF(H2F2)     HC1          HBr        HI 

Compounds  with  Oxygen. 

C12O         

CIA 

C1O2         


Compounds  ivith  Hydrogen  and  Oxygen. 

HC10  HBrO 

HC102  — 

HC103  HBrO3     HIO3 

HC104  HI04(H5I06) 

Compounds  with  One  Another. 

ClBr         101,  IBr 
IC13,IF6 

Relative  Affinities  of  the  Elements  of  the  Chlorine 
Group. — The  difference  between  the  affinities  of  these 
elements,  which  has  already  been  commented  upon,  is 
illustrated  in  a  number  of  ways.  From  iodides,  chlorine 
and  bromine  set  iodine  free  ;  and  from  bromides,  chlorine 
sets  bromine  free.  When  chlorine  is  added  to  a  solution 
containing  a  bromide  and  an  iodide,  it  first  sets  the 
iodine  free,  and  forms  the  corresponding  chloride.  Thus, 
in  the  case  of  potassium  iodide  : 

2KI  +  01,  =  2KC1  +  I2. 

After  this  reaction  is  complete,  the  chlorine  acts  upon 
the  water,  decomposing  it,  oxidizing  the  iodine  to  iodic 
acid  (which  see).  The  solution  is  then  colorless.  After 
all  the  iodine  is  converted  into  iodic  acid  the  bromine 
is  liberated  and  colors  the  solution  yellowish  red. 


1.82  INORGANIC  CHEMISTRY. 

Again,  as  we  shall  see,  there  are  some  oxidizing  agents 
which  decompose  iodides  but  which  do  not  decompose 
bromides  and  chlorides,  and  others  which  decompose 
chlorides  but  do  not  decompose  bromides  and  iodides. 

FAMILY  VII,  GROUP  A — MANGANESE. 

There  is  one  element  which  belongs  in  the  same  family 
as  those  which  have  just  been  treated,  and  resembles 
them  in  some  respects ;  but  at  the  same  time  it  differs 
from  them  quite  markedly  in  other  respects.  This  is 
manganese.  It  acts  in  fact  in  two  different  ways,  and  is 
one  of  those  elements,  already  referred  to,  which  are 
both  acid-forming  and  base-forming.  Some  of  its  com- 
pounds with  hydrogen  and  oxygen  are  distinctly  acid, 
others  are  distinctly  basic.  So  far  as  it  acts  like  the 
members  of  the  chlorine  family  a  brief  reference  to  it 
here  is  desirable.  On  the  other  hand,  it  will  be  dealt 
with  chiefly  in  connection  with  those  base-forming 
elements  which  it  most  resembles,  as,  for  example, 
iron. 

Manganese  occurs  in  nature  principally  in  the  form  of 
pyrolusite  or  manganese  dioxide,  MnO2,  also  known  as 
the  black  oxide  of  manganese.  It  forms  with  oxygen 
compounds  of  the  following  formulas :  MnO,  Mn2O3, 
Mn3O4,  MnO2,  and  Mn2O7.  When  a  compound  of  man- 
ganese is  subjected  to  the  influence  of  powerful  oxidizing 
agents  in  the  presence  of  an  alkali  it  is  converted  into  a 
salt  of  manganic  acid,  H2MnO4,  which  in  its  composition 
resembles  sulphuric  acid.  If  the  salt  of  manganic  acid 
thus  obtained  is  dissolved  in  water  it  undergoes  partial 
decomposition,  which  is  complete  if  the  solution  is  boiled, 
or  if  carbon  dioxide  is  passed  through  it.  The  change 
consists  in  the  transformation  of  manganic  acid  into  per- 
manganic acid,  HMnO4 : 

3H2MnO4  =  2HMnO4  +  MnO2  +  2H2O  ;  or 
3K2MnO4  +  2H20  =  2KMnO4  +  MnO,  +  4KOH. 

Permanganic  acid,  HMnO4,  is  a  compound  which  in 
many  respects  resembles  perchloric  acid.  It  can  be  ob- 


MANGANESE.  183 

tained  in  water  solution  by  decomposition  of  certain 
of  its  salts,  but  like  perchloric  acid  it  is  easily  decom- 
posed. In  consequence  of  the  ease  with  which  it  gives 
up  oxygen  it  is  a  good  oxidizing  agent,  and  is  extensively 
used  in  the  laboratory  in  this  capacity.  It  is  employed  in 
the  form  of  the  potassium  salt,  potassium  permanganate, 
KMnO4,  which  will  receive  special  attention  under  the 
head  of  Manganese  Compounds.  In  order,  however,  to 
make  clear  the  difference  in  conduct  between  perchloric 
and  permanganic  acids  a  few  characteristic  facts  will  be 
mentioned  here.  The  conduct  of  permanganic  acid  and 
of  potassium  permanganate  will  be  understood,  if  it  is 
borne  in  mind  that  in  the  presence  of  substances  of 
strongly  acid  character  manganese  tends  to  act  as  a  base- 
forming  element,  and  in  this  capacity  to  form  salts  with 
the  acids.  Thus  in  presence  of  sulphuric  acid  potassium 
permanganate  forms  potassium  sulphate,  manganous 
sulphate,  and  oxygen,  if  there  is  anything  present  which 
has  the  power  to  take  up  oxygen.  In  the  salts  in  which 
it  plays  the  part  of  a  metal  manganese  is  generally  biva- 
lent. With  hydrochloric  acid,  as  we  have  already  seen 
in  studying  the  action  of  hydrochloric  acid  upon  manga- 
nese dioxide,  it  forms  the  chloride  MnCl2.  When  now 
potassium  permanganate  is  treated  with  hydrochloric  acid 
it  is  decomposed  according  to  the  following  equation : 

•2KMnO4  +  6HC1  =  2KC1  +  2MnCl2  +  3H2O  +  5O. 

Similarly,  with  sulphuric  acid  manganous  sulphate  is 
formed,  thus  : 

2KMnO4  +  3H2S04  =  K2SO4  +  2MnSO4  +  3H2O  +  5O. 

Such  reactions  do  not  take  place  with  perchloric  acid,  as 
chlorine  is  entirely  lacking  in  the  power  to  enter  into 
acids  in  the  place  of  the  hydrogen  and  form  salts. 

Manganese  forms  some  other  acids  besides  perman- 
ganic acid,  but  they  exhibit  little  or  no  analogy  with 
compounds  of  chlorine,  and  their  study  will  therefore  be 
postponed  until  manganese  is  taken  up.  The  point  of 


184  INORGANIC  CHEMISTRY. 

chief  interest  to  be  noted  here  is  that  this  element  is  un* 
mistakably  like  chlorine  in  its  highest  oxygen  com- 
pounds, but  entirely  different  from  it  in  most  of  its  com- 
pounds. The  compound  manganese  heptoxide,  Mn2O7, 
stands  in  the  relation  of  an  anhydride  to  permanganic 
acid.  In  water  solution  it  passes  over  into  the  acid : 

Mn,O7  +  H2O  =  2HMnO4. 

It  is  formed  by  treating  potassium  permanganate  with 
the  most  concentrated  sulphuric  acid : 

2KMn04  +  H2S04  =  K2SO4  +  Mn2O7  +  H2O. 

It  is  extremely  unstable,  giving  up  oxygen  readily.  In 
contact  with  organic  substances  or  other  substances 
which  have  the  power  to  take  up  oxygen  it  decomposes 
so  rapidly  as  frequently  to  lead  to  explosions.  It  is  of 
interest  to  note  that  this  is  the  only  oxide  of  Family  YII 
in  which  the  maximum  valence  of  7  is  shown. 

Judging  by  analogy,  it  seems  probable  that  the  consti- 
tution of  permanganic  acid  is  like  that  of  periodic  and 
perchloric  acids,  and  is  represented  by  the  formula 

O 

ll 

O=Mn-0-H,  in  which  the  manganese  is  septivalent. 
I) 
O 


CHAPTER  XIII. 

THE  ELEMENTS  OF  FAMILY  VI,  GROUP  B : 
SULPHUR— SELENIUM— TELLURIUM. 

Introductory. — The  elements  of  this  group  bear  to  oxy- 
gen a  relation  somewhat  similar  to  that  which  the  ele- 
ments of  Group  B,  Family  VII,  bear  to  fluorine.  The 
three  members  sulphur,  selenium,  and  tellurium  resem- 
ble one  another  fully  as  strikingly  as  chlorine,  bromine, 
and  iodine  do.  Their  compounds  bear  a  general  resem- 
blance, to  those  of  oxygen,  and  yet  they  form  very  char- 
acteristic compounds  with  oxygen,  while  oxygen  forms  no 
analogous  compounds  with  any  of  them.  Just  as  iodine 
forms  a  compound  with  fluorine  of  the  formula  IF5,  but 
fluorine  does  not  form  with  iodine  a  compound  FI5,  so 
sulphur,  selenium,  and  tellurium  form  with  oxygen  the 
-compounds  SO3,  SeO3,  and  TeO3,  while  oxygen  does  not 
form  analogous  compounds  with  these  other  elements. 
The  valence  of  the  elements  of  this  group  towards  hydro- 
gen is  2,  as  shown  in  the  compounds  H2O,  H2S,  H2Se,  and 
H3Te.  Of  oxygen  and  sulphur  there  are  other  hydrogen 
compounds,  as  hydrogen  dioxide,  H2O2,  and  an  analogous 
compound  of  sulphur,  but  there  is  great  uncertainty  in 
regard  to  the  constitution  of  the  latter  compound,  and 
practically  nothing  is  known  in  regard  to  the  valence  of 
sulphur  in  it.  In  general  the  hydrogen  valence  is  con- 
stant. Towards  the  members  of  the  chlorine  group  the 
valence  varies  from  2  to  6.  Oxygen  never  exhibits  a 
higher  valence  than  2  towards  chlorine  and  its  analogues. 
The  compound  OC12  illustrates  the  bivalence  of  oxygen 
towards  chlorine.  Sulphur  forms  with  chlorine  the  com- 
pounds S2C12  and  SC12,  which  are  analogous  to  the  hy- 
drogen compounds  H2S2  and  H2S,  and  in  both  of  them 
the  sulphur  is  probably  bivalent.  It  also  forms  the  com- 
pound SC14,  in  which  it  is  quadrivalent.  With  iodine  it 

(185) 


186  INOEGANIC  CHEMISTRY. 

is  said  to  form  the  unstable  compounds  S2I2  and  SI6.  Se- 
lenium and  tellurium  form  similar  compounds,  and,  in 
general,  the  stability  of  the  compounds  of  the  members 
of  the  sulphur  group  with  the  members  of  the  chlorine 
group  increases  in  the  order  sulphur,  selenium,  tellurium. 
Sexivalence  of  these  elements  towards  members  of  the 
chlorine  group  is  rare,  being  shown  only  in  the  com- 
pound SI8,  and  even  this  is  doubtful. 

Towards  oxygen  the  three  elements  of  the  sulphur 
group  are  quadrivalent  and  sexivalent,  as  seen  in  the 
compounds  SO2,  SeO2,  TeOa,  and  SO3,  SeO3,  and  TeO3.  Of 
course,  it  is  possible  that  in  these  oxygen  compound  o 
the  elements  are  bivalent.  Thus,  sulphur  dioxide,  SO2, 

may  be   represented   by  the  formula  S^  I  ,     in    which 

both  the  oxygen  and  the  sulphur  appear  as  bivalent ;  and, 
in  a  similar  way,  the  trioxide  may  be  represented  by  the 

O 
formula  S         O ;  but  the  only  reason  for  doing  this  is 

O 

the  desire  to  represent  sulphur  as  always  bivalent.  The 
existence  of  the  compounds  SC14,  SeCl4,  and  SI6  cannot 
be  explained,  however,  on  the  assumption  that  sulphur 
is  bivalent,  and  the  simplest  view  which  can  be  taken  of 
the  matter  is  that  the  members  of  the  sulphur  group  are 
in  general  bivalent  towards  hydrogen  ;  bivalent,  quadri- 
valent, and,  exceptionally,  sexivalent  towards  the  mem- 
bers of  the  chlorine  group  ;  and  quadrivalent  and  sexi- 
valent towards  oxygen.  Towards  hydroxyl  the  valence 
of  the  members  of  the  sulphur  group  appears  to  vary 
from  4  to  6.  The  quadrivalence  is  shown  in  the  com- 
pound hydrosulphurous  acid,  H2SO2,  which  probably  has 
H 

the  constitution  O=S-O-H  ;    the  sexi valence  is  seen  in 

O 

II 
sulphurous  acid,  H2SO3,  or  H-S-O-H,  and  in  sulphuric 

II 
O 


SULPHUR.  187 

acid,    S(OH)6,    or    in    the    ordinary    form    H2S04,   or 
O 

H-O-S-O-H. 

II 
O 

Of  the  three  elements  of  this  group  sulphur  occurs  in 
greatest  abundance  in  nature,  selenium  next  in  order, 
and  finally  tellurium.  Just  as  bromine  frequently 
accompanies  chlorine,  so  selenium  frequently  accompa- 
nies sulphur,  but  it  is  always  present  in  much  smaller 
quantity  than  sulphur.  Tellurium  occurs  in  very  small 
quantity  relatively,  and  not  uncommonly  in  combina- 
tion with  valuable  metals  like  gold  and  silver.  Large 
quantities  of  sulphur  are  found  in  the  native  or  uncom- 
bined  condition.  Only  extremely  small  quantities  of 
selenium  and  tellurium  are  found  native. 


SULPHTJK,  S  (At.  Wt.  31.83). 

Occurrence. — The  principal  deposits  of  native  sulphur 
are  found  in  Sicily,  Italy,  and  Spain.  In  California  also 
there  is  a  considerable  deposit.  In  general,  sulphur  is 
likely  to  be  found  near  dying  or  extinct  volcanoes. 
Sulphur  occurs  in  nature,  further,  in  the  form  of  the 
hydrogen  compound,  hydrogen  sulphide,  H2S,  issuing 
from  the  earth  in  volcanic  regions,  and  in  solution  in 
some  natural  waters,  known  as  "sulphur  waters."  The 
oxide  SO2  is  likewise  found  issuing  from  the  earth  in 
volcanic  regions.  Compounds  of  sulphur  with  metallic 
elements,  as  with  iron,  copper,  lead,  zinc,  are  very  abun- 
dant. Such  compounds  are  iron  pyrites,  FeS2 ;  copper 
pyrites,  CuFeS2 ;  galenite,  PbS ;  and  zinc  blende,  ZnS. 
Some  sulphates  are  widely  distributed  and  occur  in 
large  quantities ;  for  example,  gypsum  or  calcium  sul- 
phate, CaSO4  +  2H2O  ;  barium  sulphate,  or  heavy  spar, 
BaSO4 ;  lead  sulphate,  PbSO4.  Finally,  sulphur  occurs 
in  a  few  animal  and  vegetable  products  in  combination 
with  carbon,  hydrogen,  and,  generally,  with  nitrogen. 

Extraction  of  Sulphur  from  its  Ores. — By  far  the  largest 
quantity  of  sulphur  found  in  the  market  is  taken  from 


188  INORGANIC  CHEMISTRY. 

the  mines  in  Sicily.  Of  these  mines  there  are  between 
250  and  300.  When  taken  from  the  mines  it  is  mixed 
with  many  earthy  substances,  from  which  it  must  be 
separated.  The  separation  is  generally  effected  in  Sicily 
by  piling  the  ore  so  as  to  leave  passages  for  air,  covering 
the  piles  with  earthy  matter  to  prevent  free  access  of  air 
and  then  setting  fire  to  them.  A  part  of  the  sulphur 
burns,  and  this  causes  the  rest  of  the  sulphur  to  melt. 
The  molten  sulphur  runs  down  to  the  bottom  of  the 
pile,  and  by  a  proper  arrangement  it  is  drawn  off  from 
time  to  time.  If  the  pile  of  ore  were  not  covered  up, 
and  the  air  were  allowed  to  enter  freely,  the  sulphur 
would  burn  up,  and  be  converted  into  the  gas  sulphur 
dioxide.  The  "  crude  brimstone  "  obtained  in  the  manner 
described  is  afterwards  refined  by  distillation,  and  it  is 
this  refined  or  distilled  sulphur  which  is  met  with  in  the 
market  under  the  names  "  roll  brimstone  "  and  "  flowers 
of  sulphur." 

The  distillation  is  carried  on  in  retorts  made  of  earth- 
enware, and  these  are  connected  with  large  chambers  of 
brick-work.  When  the  vapor  of  sulphur  first  comes  into 
the  condensing- chamber  it  is  suddenly  cooled,  and  hence 
deposited  in  the  form  of  a  fine  powder.  This  is  what  is 
called  "  flowers  of  sulphur."  After  the  distillation  has 
continued  for  some  time  the  vapor  condenses  in  the  form 
of  a  liquid,  which  collects  at  the  bottom  of  the  chamber. 
This  is  drawn  off  into  slightly  conical  wooden  moulds, 
and  takes  the  form  of  "roll  brimstone"  or  "stick  sul- 
phur." 

Some  sulphur  is  obtained  from  iron  pyrites  by  heating 
in  closed  vessels.  The  change  which  takes  place  on 
heating  iron  pyrites  is  perfectly  analogous  to  that  which 
takes  place  on  heating  manganese  dioxide  in  preparing 
oxygen.  A  sulphur  compound  of  the  formula  Fe3S4  and 
free  sulphur  are  formed  in  the  former  case,  as  the  com- 
pound of  manganese  and  oxygen,  Mn3O4,  and  free  oxygen 
are  formed  in  the  latter  : 

3FeS2    =  Fe3S4   +S2; 
3MnO2  =  Mn3O4  +  O2. 


PROPERTIES  OF  SULPHUR.  189 

Properties. — Sulphur  is  a  yellow,  brittle  substance 
which  at  —  50°  is  almost  colorless.  It  melts  at  114.5°, 
forming  a  thin,  straw-colored  liquid.  When  heated  to 
a  higher  temperature  it  becomes  darker  and  darker  in 
color,  and  at  200°  to  250°  it  is  so  viscid  that  the  vessel 
in  which  it  is  contained  may  be  turned  upside  down  with- 
out danger  of  its  running  out.  Finally,  at  448.4°  it  boils, 
and  is  then  converted  into  a  brownish-yellow  vapor. 
When  molten  sulphur  solidifies,  or  when  it  is  deposited 
from  a  solution,  it  takes  the  form  of  crystals.  But, 
strange  to  say,  the  crystals  formed  from  molten  sulphur 
are  entirely  different  from  those  deposited  from  cold 
solutions  of  sulphur.  The  latter  belong  to  the  rhombic 
system.  They  are  octahedrons  with  a  rhombic  base. 
And  this  is  the  form  of  the  sulphur  found  in  nature.  The 
former  are  honey-yellow  needles.  An  examination  of 
these  needles  shows  that  the  angles  which  their  faces 
form  with  one  another  are  not  the  same  as  the  angles 
formed  by  the  faces  of  the  octahedrons,  and  that  they 
belong  to  an  entirely  different  system — the  monoclinic. 

The  rhombic  crystals  of  sulphur  can  be  made  by  dis- 
solving "  roll  brimstone"  in  carbon  disulphide  and  allow- 
ing the  solution  to  stand.  When  the  liquid  has  suffi- 
ciently evaporated,  the  sulphur  will  appear  in  larger  or 
smaller  rhombic  crystals,  according  to  the  conditions. 
A  comparison  of  the  crystals  thus  obtained  with  natural 
crystals  will  show  that  the  two  have  identical  or  very 
similar  forms.  The  formation  of  the  needles  or  mono- 
clinic  crystals  may  be  shown  by  melting  a  considerable 
quantity,  say  a  pound  or  two,  of  roll  brimstone  in  a  sand 
•or  Hessian  crucible,  and  allowing  the  liquid  mass  to  cool 
slowly.  When  a  thin  crust  has  formed  on  the  surface 
this  should  be  perforated,  and  the  remainder  of  the 
liquid  sulphur  poured  off.  The  crucible  will  then  be 
found  lined  with  long,  dark  yellow,  lustrous  needles 
which  do  not  look  at  all  like  those  obtained  from  the 
solution  in  carbon  disulphide.  If  the  monoclinic  needles 
a-re  allowed  to  lie  unmolested  they  gradually  undergo 
change  spontaneously.  They  lose  their  lustre  and  be- 
come lighter  in  color ;  and  now,  if  examined  carefully,  they 


190  INORGANIC  CHEMISTRY. 

are  found  to  consist  of  minute  crystals  like  those  found 
in  nature.  They  have  changed  to  the  rhombic  form.  It 
is  evident,  therefore,  that  the  arrangement  of  the  parti- 
cles in  the  monoclinic  crystals  of  sulphur  is  not  a  stable 
one.  The  change  is  accompanied  by  a  considerable  evo- 
lution of  heat. 

Substances  which  crystallize  in  two  distinct  forms  are 
called  dimorphous.  We  shall  see  that  carbon  also  crys- 
tallizes in  two  different  forms,  and  that  this  kind  of  phe- 
nomenon is  met  with  not  unfrequently  among  chemical 
compounds.  The  difference  between  the  two  varieties 
of  sulphur  suggests  that  observed  between  the  two 
forms  of  oxygen.  Whether  the  explanation  is  the  same 
in  the  two  cases  is  doubtful.  It  appears  more  probable 
that  the  difference  in  the  former  case  is  due  to  different 
arrangements  of  the  molecules  in  the  crystals,  rather 
than  to  different  arrangements  of  the  atoms  in  the  mole- 
cules. The  chemical  properties  of  the  two  varieties  are 
practically  identical.  This  could  hardly  be  the  case  if 
the  number  of  atoms  in  the  molecule  were  different  in 
the  two  cases. 

Besides  the  two  forms  mentioned  sulphur  can  also  be 
obtained  in  the  amorphous,  or  uncrystallized,  condition. 
If  molten  sulphur  is  quickly  cooled  under  water  it  re- 
mains for  some  time  soft  and  dough-like,  and  while  in 
this  condition  it  is  amorphous.  If  allowed  to  stand  it 
gradually  becomes  hard  and  brittle. 

When  separated  from  certain  compounds  which  are 
in  solution  in  water,  the  sulphur  is  in  a  very  finely 
divided  condition,  and  gives  the  liquid  the  appearance  of 
milk.  This  is  seen  on  adding  hydrochloric  acid  to  a 
solution  of  sodium  thiosulphate,  or  hyposulphite,  as  it 
is  generally  called.  This  substance  has  the  formula 
Na2S2O3,  and  the  reactions  which  take  place  between  it 
and  hydrochloric  acid  are  these : 

Na2S203  +  2HC1  =  H2S203  +  2NaCl ; 
H,S203  =  SO,  +  H20  +  S. 

On  treating  certain  varieties  of  sulphur  with  carbon 
disulphide  they  are  found  to  dissolve  completely.  This 


PROPERTIES  OF  SULPHUR.  191 

is  true,  for  example,  of  the  natural  crystals  and  of  those 
made  artificially  by  deposition  from  a  solution  in  carbon 
disulphide.  On  the  other  hand,  sulphur  in  the  form  of 
"  flowers  of  sulphur  "  is  only  partly  soluble  in  the  liquid. 
There  are  therefore  two  forms  of  sulphur  to  be  dis- 
tinguished between,  the  soluble  and  the  insoluble.  The 
cause  of  the  difference  between  these  modifications  is 
not  known.  "  Stick  sulphur  "  is  mostly  soluble,  while 
in  the  "  flowers  of  sulphur  "  there  is  at  times  a  consid- 
erable percentage  of  the  insoluble  variety.  Sulphur  is 
insoluble  in  water,  and  slightly  soluble  in  alcohol  and 
ether. 

Sulphur  is  a  much  less  active  element  chemically  than 
the  members  of  the  chlorine  group,  and  also  less  active 
than  oxygen.  Generally  speaking,  however,  it  conducts 
itself  like  oxygen.  It  combines  directly  though  not 
easily  with  hydrogen,  and  it  combines  readily  with  most 
metals,  forming  compounds  called  sulphides  which,  so  far 
as  their  composition  is  concerned,  are  analogous  to  the 
oxides.  Thus  when  heated  together  with  iron,  copper, 
or  lead,  combination  takes  place  readily  with  evolution 
of  heat  and  light.  In  its  chemical  conduct  it  differs 
markedly  from  the  members  of  the  chlorine  group. 
When  heated  to  a  sufficiently  high  temperature  in  the  air 
or  in  oxygen,  sulphur  forms  the  compound  sulphur 
dioxide,  SO2,  and,  under  certain  conditions  which  will  be 
described  farther  on,  this  combines  with  more  oxygen 
to  form  the  trioxide  SO3.  Further,  it  combines  with 
nearly  all  the  acid-forming  elements  if  heated  with  them 
to  a  sufficiently  high  temperature.  It  combines  with 
most  other  elements  less  readily  than  oxygen,  and  it 
forms  less  stable  compounds.  Thus,  its  compound  with 
hydrogen  is  decomposed  very  much  more  readily  into  its 
elements  than  water  is ;  and  the  sulphides  or  its  com- 
pounds with  metals  are  decomposed  when  they  are 
heated  in  oxygen,  the  oxygen  displacing  the  sulphur, 
much  as  chlorine  displaces  bromine ;  though  there  is  a 
difference  between  the  two  cases  to  be  found  in  the  fact 
that  sulphur  itself  unites  readily  with  oxygen,  and  this 
facilitates  the  decomposition  of  the  sulphides  by  the 
action  of  oxygen. 


192  INORGANIC  CHEMISTRY. 

When  treated  with  powerful  oxidizing  agents  sulphur 
is  converted  into  sulphuric  acid.  Thus  the  action  of 
concentrated  nitric  acid  takes  place  in  the  main  accord- 
ing to  the  equation 

2HNO3  +  S  =  H2SO4  +  2NO. 

The  action  of  sulphur  upon  the  so-called  caustic  al- 
kalies, sodium  and  potassium  hydroxides,  is  somewhat 
like  that  of  chlorine,  bromine,  and  iodine.  It  will  be 
remembered  that  with  potassium  hydroxide  chlorine 
forms  potassium  chloride  and  potassium  chlorate  or 
hypochlorite,  according  to  the  concentration  and  tem- 
perature of  the  solution.  When  sulphur  acts  upon 
sodium  hydroxide  the  sulphide  is  formed,  but  oxygen  is 
thus  rendered  available  and  some  of  it  combines  with  the 
sulphide,  and  sulphur  also  combines  with  a  part  of  the 
sulphide.  The  principal  reactions  involved  are  : 


(1)  2NaOH  +  S 

(2)  Na2S      +  4S 

(3)  Na2S      +  S    -f  30  =  Na2S2O3. 

Expressing  these  reactions  in  one  equation  we  have 
GNaOH  +  12S  =  2Na2S&  +  Na2S2O3  +  3H2O. 


The  action  is  of  the  same  general  character  as  that 
which  takes  place  in  the  case  of  chlorine,  but  differs  from 
it  in  the  fact  that  sodium  sulphide,  Na2S,  has  the  power 
to  take  up  sulphur,  and  also  to  take  up  sulphur  and 
oxygen. 

The  specific  gravity  of  the  vapor  of  sulphur  varies  with 
the  temperature,  and  is  such  as  to  indicate  that  at  tem- 
peratures not  far  above  the  boiling  point  the  molecule 
consists  of  eight  atoms,  while  at  the  temperature  800° 
and  higher  the  molecule  consists  of  two  atoms.  This 
suggests  the  question  whether  the  molecule  of  sulphur 
in  the  solid  form  may  not  be  more  complex  than  the 
condition  represented  by  the  symbol  Sb.  We  have  no 
means  of  answering  this  question  with  any  certainty  at 
present. 


COMPOUNDS  OF  SULPHUR   WITH  HYDROGEN.      193 

(Jses  of  Sulphur. — Enormous  quantities  of  sulphur  are 
used  in  the  manufacture  of  sulphuric  acid,  and  of  gun- 
powder. It  is  also  used  in  the  manufacture  of  fire-works 
of  various  kinds.  Burning  sulphur  gives  sulphur  dioxide, 
which  is  extensively  employed  for  bleaching  wool,  silk, 
straw,  etc.  When  caoutchouc  is  thoroughly  mixed  with 
sulphur  or  some  sulphur  compound  it  becomes  vulcan- 
ized. 

COMPOUNDS -OF  SULPHUE  WITH  HYDEOGEN. 

The  principal  compound  of  sulphur  and  hydrogen  is 
analogous  in  composition  to  water.  It  is  known  as  hy- 
drogen sulphide  or  sulphuretted  hydrogen,  and  has  the 
formula  H2S.  Besides  this  there  is  at  least  one  other 
compound  which  contains  a  larger  proportion  of  sulphur, 
and  probably  has  the  composition  H2S2.  There  are 
reasons  for  supposing,  further,  that  still  more  complex 
compounds  can  exist,  but  owing  to  their  instability  it  is 
impossible  to  isolate  them  in  pure  condition  and  study 
them. 

Hydrogen  Sulphide,  Sulphuretted  Hydrogen,  H2S. — 
When  hydrogen  is  passed  over  highly  heated  sulphur 
the  two  elements  combine  to  form  hydrogen  sulphide. 
The  action  is,  however,  quite  incomplete  and  is  not  to  be 
compared  with  that  which  takes  place  when  hydrogen  and 
oxygen  are  heated  together.  This  compound  of  sulphur 
and  hydrogen  occurs  in  nature  in  solution  in  the  so-called 
"  sulphur  waters,"  which  are  met  with  in  many  parts  of 
the  world.  It  is  formed  by  heating  organic  substances 
which  contain  sulphur,  just  as  water  is  formed  by  heat- 
ing organic  substances  which  contain  oxygen.  It  is 
formed,  further,  by  decomposition  of  organic  substances 
which  contain  sulphur,  as,  for  example,  the  albumen  of 
eggs.  The  odor  of  rotten  eggs  is  partly  due  to  the  for- 
mation of  hydrogen  sulphide.  It  is  formed  by  the  action 
of  acids  upon  sulphides  or  hydrosulphides,  just  as  water 
is  formed  by  the  action  of  acids  upon  oxides  or  hydrox- 
ides (see  p.  132).  Thus  hydrochloric  acid  and  ferrous 
sulphide,  FeS,  give  ferrous  chloride,  FeCl2,  and  hydro- 
gen sulphide : 


194  INORGANIC  CHEMISTRY. 

FeS  +  2HC1  =  FeCla  +  H2S ; 

just  as  ferrous  oxide,  FeO,  and  hydrochloric  acid  give 
ferrous  chloride  and  water  : 

FeO  +  2HC1  =  FeCl2  +  H2O. 

So  also  potassium  hydroxide  and  potassium  hydrosul- 
phide  act  in  the  same  way,  as  has  been  pointed  out : 

KSH  +  HC1  =  KC1  +  H2S  ; 
KOH  +  HC1  =  KC1  +'H20. 

It  is  generally  formed  by  the  action  of  nascent  hydro- 
gen upon  sulphur  compounds.  Thus,  it  has  been  shown 
that  the  hydrogen  from  hydriodic  acid  has  the  power  to 
reduce  sulphuric  acid  to  hydrogen  sulphide : 

H2S04  +  8HI  =  H2S  +  4H20  +  81. 

In  the  laboratory,  where  the  gas  is  extensively  used, 
it  is  generally  prepared  from  ferrous  sulphide,  FeS,  and 
dilute  sulphuric  acid,  which  are  simply  brought  together 
at  the  ordinary  temperature  in  a  flask  such  as  is  used  in 
making  hydrogen.  The  reaction  is  like  that  which  takes 
place  between  ferrous  sulphide  and  hydrochloric  acid. 
It  is  represented  by  this  equation : 

FeS  +  H2S04  =  FeS04  +  H2S. 

Properties. — Hydrogen  sulphide  is  a  colorless,  trans- 
parent gas  of  the  specific  gravity  1.178.  It  has  an  ex- 
tremely disagreeable  odor,  somewhat  suggestive  of  that 
of  rotten  eggs.  It  is  poisonous,  even  small  quantities 
causing  headache,  vertigo,  nausea,  and  other  bad  symp- 
toms. It  burns  with  a  blue  flame,  forming  water  and 
sulphur  dioxide : 

H2S  +  3O  =  H20  +  S02. 

If,  however,  the  air  has  not  free  access,  as  when  the  gas 
is  burned  in  a  cylinder  open  at  one  end,  only  a  part  of 
the  sulphur  is  burned,  the  rest  being  deposited  upon 
the  walls  of  the  vessel,  while  the  hydrogen  burns.  The 
gas  is  soluble  in  water,  about  three  volumes  being  taken 


PROPERTIES  OF  HYDROGEN  SULPHIDE.  195 

up  at  ordinary  temperatures.  This  solution  is  used  to 
some  extent  in  the  laboratory  instead  of  the  gas,  but, 
owing  to  the  fact  that  it  readily  undergoes  change  in 
consequence  of  the  action  of  the  oxygen  of  the  air,  it  is 
not  as  valuable  as  it  would  be  if  it  were  more  stable. 
The  change  consists  simply  in  the  oxidation  of  the  hy- 
drogen and  the  separation  of  the  sulphur.  If  a  bottle 
containing  a  solution  of  hydrogen  sulphide  is  allowed 
to  stand  for  a  few  days,  particularly  if  it  is  opened  from 
time  to  time,  the  odor  of  the  gas  will  disappear  and  a 
deposit  of  sulphur  will  be  noticed  in  the  bottle.  The 
liquid  is  then  nothing  but  water.  When  the  solution  is 
boiled  it  loses  all  the  gas  contained  in  it. 

Hydrogen  sulphide  is  easily  decomposed  into  its  ele- 
ments. It  requires  a  temperature  of  only  a  little  above 
400°  to  effect  direct  decomposition.  In  consequence  of 
this  instability  it  causes  a  number  of  changes  which  the 
analogous  compound  water  cannot  effect.  The  relations 
here  are  similar  to  those  which  exist  between  hydro- 
chloric and  hydriodic  acids.  Hydrochloric  acid  is  very 
stable,  while  hydriodic  acid  breaks  down  readily  into 
hydrogen  and  iodine.  Therefore  hydriodic  acid,  as  we 
have  seen,  acts  as  a  reducing  agent,  while  hydrochloric 
acid  does  not.  So,  also,  hydrogen  sulphide  acts  as  a 
reducing  agent.  Thus,  if  it  be  passed  into  concentrated 
sulphuric  acid  this  reaction  takes  place : 

H2S04  +  H2S  =  2H20  +  S  +  S02. 

The  action  is  to  be  traced  to  the  decomposition  of  the 
hydrogen  sulphide  into  hydrogen  and  free  sulphur,  the 
hydrogen  then  acting  upon  the  sulphuric  acid  thus : 

H2S04  +  2H  =  2H20  +  S02. 

With  hydriodic  acid  the  reduction  may  go  farther,  as 
has  been  seen ;  with  hydrobromic  acid,  however,  the 
action  takes  place  practically  in  the  same  way  as  with 
hydrogen  sulphide. 

Chlorine,    bromine,    and   iodine   act  upon   hydrogen 


196  INORGANIC  CHEMISTRY. 

sulphide,  setting  the  sulphur  free  and  uniting  with  the 
hydrogen : 

H,S  +  01,  =  2HC1  +  S. 

This  reaction  suggests  the  decomposition  of  water  by 
chlorine,  but  what  a  difference  there  is  between  the  two 
cases !  Chlorine  decomposes  water  very  slowly  and 
only  under  the  influence  of  the  direct  sunlight,  while  it 
decomposes  hydrogen  sulphide  completely  and  instantly. 
Similarly,  hydrogen  sulphide  has  the  power  to  abstract 
chlorine  from  some  of  its  compounds,  as,  for  example, 
from  ferric  chloride,  FeCl3.  When  this  is  treated  with 
nascent  hydrogen  from  any  source,  it  is  reduced  to  fer- 
rous chloride,  FeCl2,  thus : 

FeCl3  +  H  =  FeCl2  +  HC1. 

So,  also,  when  it  is  treated  with  hydrogen  sulphide  it 
is  reduced  in  the  same  way  in  consequence  of  the  action 
of  the  hydrogen : 

2FeCl3  +  H2S  =•  2FeCl2  +  2HC1  +  S. 

The  instability  of  hydrogen  sulphide  is  further  shown 
by  the  ease  with  which  it  is  decomposed  by  metals  with 
liberation  of  hydrogen  and  formation  of  sulphides.  It 
will  be  remembered  that  at  high  temperatures  several 
metals  decompose  water,  but  that  at  ordinary  tempera- 
tures only  a  few  decompose  it  easily.  Hydrogen  sul- 
phide acts  much  more  readily  ;  a  number  of  metals  which 
do  not  act  upon  water  even  at  high  temperatures,  as  sil- 
ver, gold,  and  mercury,  decompose  this  gas  at  ordinary 
temperatures. 

Hydrogen  sulphide  acts  upon  metallic  oxides,  convert- 
ing them  into  sulphides,  as  for  example : 

CuO  +  H2S  =  CuS  +  H2O. 

Action  of  Hydrogen  Sulphide  upon  Solutions  of  Salts- 
Use  in  Chemical  Analysis. — Hydrogen  sulphide  is  exten- 
sively used  in  every  chemical  laboratory  as  a  reagent  in 
chemical  analysis.  In  order  that  its  action  may  be 
understood  a  few  words  of  explanation  are  necessary. 
Sulphur,  as  we  have  seen,  has  a  strong  affinity  for  the 


HYDROGEN  SULPHIDE  IN  CHEMICAL  ANALYSIS.  197 

metallic  or  base-forming  elements,  forming  with  them 
the  sulphides.  Further,  hydrogen  sulphide  is  easily 
decomposed,  and  the  substitution  of  metals  for  the 
hydrogen  is  facilitated  by  this  fact.  If  now  a  salt  is  in 
solution  in  water  and  hydrogen  sulphide  is  passed  into 
the  solution  there  will,  of  course,  be  the  tendency  to  the 
formation  of  the  sulphide  of  the  metal  contained  in  the 
salt.  Thus,  suppose  the  salt  in  solution  is  silver  nitrate, 
AgNO3.  On  passing  hydrogen  sulphide  into  this  solu- 
tion the  silver  will  tend  to  combine  with  the  sulphur  to 
form  the  sulphide,  Ag2S.  If  this  is  formed,  hydrogen 
must  be  freed  from  the  hydrogen  sulphide,  and  this 
would  probably  take  the  place  of  the  silver  in  the  nitrate, 
forming  nitric  acid,  according  to  this  equation : 

2AgNO3  +  H2S  =  Ag2S  +  2HNO3. 

If  the  dilute  acid  thus  formed  has  the  power  to  decom- 
pose silver  sulphide  the  sulphide  will  not  be  formed ; 
but  if  it  has  not  this  power  the  sulphide  will  be  formed, 
and  it  will  be  thrown  down  or  precipitated  if  it  is  an 
insoluble  compound.  The  sulphides  of  some  metals 
are  not  decomposed  by  dilute  acids,  and  are  insoluble 
in  water.  If  hydrogen  sulphide  is  passed  through  solu- 
tions of  the  salts  of  these  metals  the  sulphides  are 
thrown  down. 

Secondly,  the  sulphides  of  some  metals  are  decom- 
posed by  dilute  acids.  Plainly,  these  cannot  be  thrown 
down  by  simply  passing  hydrogen  sulphide  through  the 
solutions  of  their  salts,  whether  they  are  soluble  or  in- 
soluble in  water.  Thus,  for  example,  the  sulphide  of 
iron,  FeS,  is  insoluble  in  water,  but  it  is  easily  decom- 
posed by  dilute  acids,  and  therefore  when  hydrogen 
sulphide  is  passed  through  a  solution  of  an  iron  salt  the 
sulphide  is  not  precipitated.  In  the  case  of  the  sulr 
phate  the  reaction  would  be 

FeS04  +  H2S  =  FeS  +  H2SO4. 

But  the  dilute  sulphuric  acid  would  decompose  the  sul- 
phide, and  the  reaction  does  not  take  place.  If,  however, 
a  soluble  sulphide  is  added  to  a  solution  of  such  a  metal, 


198  INORGANIC  CHEMISTRY. 

decomposition  takes  place.  Thus,  if,  instead  of  passing 
hydrogen  sulphide,  a  solution  of  potassium  sulphide  is 
added,  reaction  takes  place  thus : 

FeS04  -f  K2S  =  K2S04  +  FeS. 

Here  no  sulphuric  acid  is  formed,  but,  instead  of  it, 
neutral  potassium  sulphate,  which  does  not  act  upon  the 
insoluble  sulphide.  Advantage  is  taken  of  this  fact  in 
chemical  analysis,  but,  in  place  of  potassium  sulphide, 
the  analogous  compound,  ammonium  sulphide,  (NH4)2S, 
is  used.  This  acts  in  the  same  way.  Thus,  in  the  case 
above  cited  the  action  with  ammonium  sulphide  is  repre- 
sented by  this  equation : 

FeSO4  +  (NH4)aS  =  (NH4)2SO4  +  FeS. 

There  are  several  metals  which  act  towards  hydrogen 
sulphide  in  the  same  way  that  iron  does. 

Thirdly,  there  are  some  metals  whose  sulphides  are 
soluble  in  water,  and,  therefore,  if  solutions  of  their  salts 
are  treated  with  hydrogen  sulphide  or  with  ammonium 
sulphide  no  apparent  action  takes  place. 

Facts  like  those  just  referred  to  form  a  good  basis 
for  the  division  of  the  metallic  elements  into  groups  for 
purposes  of  analysis.  According  to  the  above  these  ele- 
ments can  be  divided  into  three  great  groups,  as  follows : 

I.  Metals  whose  sulphides  are  insoluble  in  water  and 
not  decomposed  by  dilute  acids.     This  is  called  the  hy- 
drogen sulphide  group.     It  includes  lead,  bismuth,  silver, 
mercury,    copper,   cadmium,  gold,   platinum,  tin,   anti- 
mony, and  arsenic. 

II.  Metals  whose  sulphides  are  insoluble  in  water  but 
are  decomposed  by  dilute  acids.     They  are  therefore 
not  precipitated  by  hydrogen  sulphide,  but  are  precipi- 
tated by  soluble  sulphides.     As  ammonium  sulphide  is 
used  for  the  purpose  of  effecting  the  precipitation  the 
group  is  known  as  the  ammonium  sulphide  group.     It  in- 
cludes iron,  nickel,  cobalt,  manganese,  thallium,   zinc, 
and  uranium.     Further,  the  two  elements  aluminium  and 
chromium  are  thrown  down  with  the  above,  not  as  sul- 


HYDROGEN  SULPHIDE  IN  CHEMICAL  ANALYSIS.    199 

phides  but  as  hydroxides,  and  they  are  therefore  in- 
cluded in  the  group. 

III.  Metals  whose  sulphides  are  soluble  in  water. 
This  group  includes  all  the  metals  not  included  in  the 
above  two. 

By  taking  advantage,  then,  of  the  properties  of  the 
sulphides  of  the  metals  they  can  be  divided  into  these 
three  groups,  and  the  detection  of  any  particular  element 
is  thus  facilitated.  If  hydrogen  sulphide  is  passed 
through  a  solution,  and  a  precipitate  formed,  we  know 
that  this  can  contain  only  those  metals  which  belong  to 
the  hydrogen  sulphide  group ;  and  so  on.  Now,  if  the 
precipitate  formed  with  hydrogen  sulphide  is  treated 
with  certain  other  reagents  other  changes  take  place, 
and  by  further  study  it  is  quite  possible,  and  indeed 
comparatively  simple,  to  determine  which  of  the  mem- 
bers of  the  group  are  present. 

One  reaction  made  use  of  in  further  examination  of 
the  hydrogen  sulphide  precipitate  is  particularly  inter- 
esting in  this  connection.  Under  the  head  of  Acids  and 
Bases  compounds  were  referred  to  which  were  called  sul- 
phur acids  and  sulphur  bases.  Corresponding  to  the 
oxygen  acid  known  as  arsenic  acid,  which  has  the  formula 
H3AsO4,  there  are  salts  which  are  plainly  derived  from 
the  corresponding  sulphur  acid  H3AsS4.  Such  salts  are 
formed  by  treating  arsenic  sulphide  with  soluble  sul- 
phides, as  for  example  ammonium  sulphide  : 

As2S5  +  3(NH4)2S  =  2(NH4)3AsS4. 

So,  too,  tin  forms  an  oxygen  acid,  H2SnO3 ;  and  salts  of 
the  corresponding  sulphur  acid,  H2SnS3,  are  formed  by 
treating  the  sulphide  of  tin  with  soluble  sulphides : 

SnS2  +  (NH4)2S  =  (NH4)2SnS3. 

Now  some  of  these  sulphur  salts  are  soluble  in  water. 
This  is  true  of  the  ammonium  salts.  So  that  when 
the  sulphides  of  metals  which  have  the  power  to  form 
salts  of  this  kind  are  treated  with  ammonium  sul- 
phide they  pass  into  solution.  Of  the  metals  of  the 


200  INORGANIC  CHEMISTRY. 

hydrogen  sulphide  group  only  arsenic,  antimony,  and 
tin  have  this  power  ;  so  that,  if  the  hydrogen  sul- 
phide precipitate  is  treated  with  ammonium  sulphide, 
arsenic,  antimony,  and  tin  sulphides  are  dissolved  if 
present,  whereas  the  other  sulphides  are  not  changed 
by  this  treatment.  Thus  a  means'is  afforded  of  subdivid- 
ing the  hydrogen  sulphide  group  into  two  sub-groups: 
(a)  Metals  whose  sulphides  are  insoluble  in  ammonium 
sulphide  ;  and  (6)  metals  whose  sulphides  are  soluble 
in  ammonium  sulphide. 

Hydrosulphid.es.  —  The  action  of  hydrogen  sulphide 
shows  that  it  belongs  to  the  class  of  acids.  When  it 
acts  upon  an  oxide  the  corresponding  sulphide  and 
water  are  formed.  But  just  as  there  are  sulphides 
which  are  derived  from  hydrogen  sulphide  by  the  re- 
placement of  both  hydrogen  atoms  by  metallic  elements, 
so  there  are  hydrosulphides  which  are  derived  from  it 
by  the  replacement  of  only  one  of  the  two  atoms  of  hy- 
drogen in  the  molecule.  The  sulphides  correspond  to 
the  oxides,  and  the  hydrosulphides  to  the  hydroxides. 
Thus  the  analogous  oxygen  and  sulphur  compounds  of 
potassium  are  : 


K20 

KOH        KSH. 

We  speak  of  sulphides  and  hydrosulphides  as  salts  of  hy- 
drogen sulphide.  In  consequence  of  this  ability  to  form 
salts  in  the  same  way  in  general  that  acids  do,  hydrogen 
sulphide  is  sometimes  called  sulphydric  acid,  and  the  salts 
of  the  formula  MSH,  sulphydrates.  The  name  sulphydrate 
is  analogous  to  hydrate,  which,  as  has  been  pointed  out, 
is  used  by  some  to  designate  the  compounds  of  the  for- 
mula MOH  or  the  hydroxides.  Between  the  acid,  hydro- 
gen sulphide,  and  the  neutral  compound,  water,  there  is  no 
fundamental  difference.  The  difference  is  simply  one  of 
degree.  In  the  present  system  of  chemistry,  which  is 
largely  an  oxygen  system,  water  is  regarded  as  the  con- 
necting link  between  acids  and  bases,  as  was  shown  on 
p.  135.  But  we  might  with  equal  right  base  our  defi- 


HYDROGEN  PERSULPHIDE.  201 

nitions  and  conceptions  of  acids  and  bases  upon  the 
conduct  of  sulphur  compounds,  and  thus  build  up  a 
sulphur  system.  In  such  a  system  hydrogen  sulphide 
would  be  the  connecting  link  between  acids  and  bases. 

Hydrogen  Persulphide,  HaSa(P). — The  sulphides  of  cer- 
tain metals,  particularly  the  so-called  alkali  metals,  so- 
dium and  potassium,  combine  with  sulphur  to  form  the 
polysulphides,  examples  of  which  are  K2S2,  K2S3,  K2S4, 
and  K2S5.  When  these  are  decomposed  with  dilute  acids 
compounds  of  hydrogen  and  sulphur  are  formed.  It 
has  thus  far,  however,  been  impossible  to  determine 
whether  more  than  one  such  compound  is  formed,  for 
the  reason  that  there  is  no  means  of  deciding  whether 
the  substances  formed  are  chemical  compounds  or  mere 
mixtures  of  sulphur  and  some  one  compound  of  sulphur 
and  hydrogen.  Hydrogen  persulphide  is  a  liquid  with  a 
very  disagreeable  odor.  Just  as  hydrogen  dioxide  de- 
composes readily  into  water  and  oxygen,  so  hydrogen 
persulphide  decomposes  readily  into  hydrogen  sulphide 
and  sulphur. 

Concerning  the  constitution  of  hydrogen  persulphide 
nothing  definite  is  known.  If  the  constitution  of  hydro- 
gen dioxide  is  H-O-O-H,  and  hydrogen  persulphide  is 
a  disulphide,  it  seems  probable  that  it  has  the  constitu- 
tion H-S-S-H,  but  there  is  no  evidence  bearing  upon 
this  point.  The  fact  that  the  sulphides  can  take  up  four 
and  only  four  atoms  of  sulphur,  just  as  they  can  take  up 
four  and  only  four  atoms  of  oxygen,  taken  together  with 
the  general  similarity  between  the  conduct  of  oxygen  and 
that  of  sulphur,  suggests  that  these  two  reactions  may 
be  of  the  same  kind : 

K2S  +  4S  =  K2SS4 ; 
K2S  +  40  =  K2SO4. 

But  for  reasons  which  will  be  more  fully  presented 
under  Sulphuric  Acid  (which  see),  it  is  generally  believed 
that  in  this  acid  two  of  the  oxygen  atoms  are  in  direct 
combination  with  sulphur  alone,  while  two  are  in  com- 
bination with  sulphur  and  hydrogen.  If  the  polysul- 


202  INORGANIC  CHEMISTRY. 

phide  has  a  similar  constitution  it  must  be  represented 

S 

II 
by  the  formula  K-S-S-S-K,  and,  further,  if  the  structure 

II 
8 

of  the  polysulphide  be  as  here  represented,  it  is  possible 
that  persulphide  of  hydrogen  has  a  similar  constitution. 
Compounds  of  Sulphur  with  Members  of  the  Chlorine 
Group. — Sulphur  combines  directly  with  chlorine  and 
forms  the  compounds  S2C12,  SC12,  and  SC14.  Of  these 
the  first  is  the  most  stable.  This  can  be  boiled  without 
undergoing  decomposition.  The  second,  sulphur  dichlo- 
ride,  SC12,  undergoes  decomposition  into  chlorine  and 
sulphur  monochloride  at  the  boiling  point ;  while  sulphur 
tetrachloride  exists  only  at  low  temperatures.  All  these 
compounds  are  decomposed  by  water,  yielding  oxygen 
compounds.  In  referring  to  the  differences  between  the 
acid-forming  and  the  base-forming  elements,  attention 
was  called  (see  p.  125)  to  the  fact  that,  in  general,  the 
oxides  of  the  base-forming  elements  are  decomposed  by 
hydrochloric  acid,  yielding  metallic  chlorides  and  water ; 
whereas  with  the  acid-forming  elements  the  reverse  is 
true,  that  is  to  say,  the  chlorides  of  the  acid-forming 
elements  are,  in  general,  decomposed  by  water  yielding 
oxides  or  hydroxides  and  hydrochloric  acid.  The  truth 
of  this  general  statement  is  illustrated  by  the  compounds 
of  sulphur  and  chlorine.  But  the  products  formed,  ex- 
cept in  the  case  of  the  tetrachloride,  are  not  strictly 
analogous  to  the  chlorides  which  are  decomposed.  Thus, 
if  in  the  monochloride  S2C12  oxygen  were  simply  substi- 
tuted for  chlorine,  the  product  would  be  S2O  ;  but  there 
is  no  compound  of  oxygen  of  this  composition,  the  sim- 
plest one  being  sulphur  dioxide,  SO2,  and  this  is  formed. 
The  excess  of  sulphur  set  free,  above  that  required  for 
the  formation  of  the  dioxide,  is  precipitated.  The  main 
part  of  the  reaction  is  represented  thus  : 

2S2C12  +  2H20  =  4HC1  +  SO2  +  38. 
The  decomposition  of  the  other  chlorides  takes  place  in 


SELENIUM.  203 

a  similar  way.     That  of  the  dichloride  is  represented 
by  the  equation 

2SC12  +  2H2O  =  SO2  +  4HC1  +  S. 

That  of  the  tetrachloride  consists  simply  in  the  direct 
replacement  of  the  chlorine  by  oxygen. 

Of  the  other  compounds  of  sulphur  with  members  of 
the  chlorine  group,  the  hexiodide,  SI6,  is  perhaps  the 
most  interesting,  as  it  shows  that  sulphur  can  be  sexiva- 
lent  towards  other  elements  than  oxygen.  This  hexio- 
dide is  an  extremely  unstable  compound.  There  is 
doubt  as  to  its  existence. 

SELENIUM,  Se  (At.  Wt.  78.42). 

Occurrence. — Selenium  occurs  only  in  small  quantity 
yn  nature.  It  was  first  found  in  the  deposit  formed  in  a 
sulphuric  acid  chamber  (see  p.  215),  and  owed  its  origin 
to  the  presence  of  small  quantities  of  selenides  in  the  sul- 
phides used  in  the  operation.  It  was  found  to  resemble 
tellurium,  and  for  that  reason  was  called  selenium,  from 
<re\rjvrj,  the  moon,  tellurium  receiving  its  name  from 
tdlus,  the  earth.  Selenium  occurs  in  a  number  of  the 
natural  sulphides,  such  as  iron  pyrites,  copper  pyrites, . 
zinc  blende,  etc. ;  and  when  these  are  treated  in  a  current 
of  air  to  decompose  them  and  convert  the  sulphur  into 
sulphur  dioxide,  the  selenium  is  also  oxidized,  and  the 
.selenium  dioxide  thus  formed  is  carried  into  the  flues 
and  other  parts  of  the  apparatus.  As  it  is  a  solid  it  is 
easily  condensed,  and  gradually  a  considerable  quantity 
collects  in  the  flues.  This  flue-dust  is  the  best  material 
from  which  to  make  selenium.  For  the  purpose  of  ex- 
traction the  dust  is  treated  with  oxidizing  agents,  and 
the  selenium  converted  either  into  selenious  acid,  H2SeO3, 
or  a  salt  of  selenic  acid,  H2SeO4.  This  is  then  reduced 
by  proper  means.  Thus  selenious  acid  is  reduced  to 
selenium  by  passing  sulphur  dioxide  through  its  solu- 
tion : 

HJ3eOa  +  2SO.  +  H,O  =  Se  4-  2HJ3O.. 


204  INORGANIC  CHEMISTRY. 

Hydrogen  sulphide  also  reduces  selenious  acid,  forming, 
however,  not  free  selenium,  but  selenium  sulphide,  SeS2, 
analogous  to  sulphur  dioxide  and  selenium  dioxide : 

H2Se03  +  2H2S  =  3H2O  +  SeS2. 

Properties. — There  are  two  modifications  of  selenium, 
corresponding  to  those  of  sulphur.  One  is  soluble  in 
carbon  disulphide,  the  other  is  not.  The  soluble  form 
is  obtained  by  reducing  selenious  acid  by  means  of  sul- 
phurous acid,  or  by  means  of  other  reducing  agents.  The 
insoluble  variety  is  obtained  by  melting  selenium,  then 
cooling  it  down  suddenly  to  210°,  and  keeping  it  at 
this  temperature  for  some  time.  When  it  solidifies  it  is 
found  to  be  no  longer  soluble  in  carbon  disulphide.  Se- 
lenium burns  as  sulphur  does,  forming  the  oxide  SeO2, 
which  has  the  odor  of  rotten  horse-radishes. 

Hydrogen  Selenide,  H2Se. — This  compound  is  made  in 
the  same  way  as  the  corresponding  compound  of  sulphur, 
by  treating  a  selenide,  as,  for  example,  ferrous  selenide, 
FeSe,  with  an  acid.  It  is  a  gas,  which  is  soluble  in 
water,  and  has  an  odor  something  like  that  of  hydrogen 
sulphide,  but  it  produces  much  more  powerful  effects 
upon  the  olfactory  nerves,  temporarily  destroying  the 
.sense  of  smell  and  causing  painful  sensations. 

The  compounds  of  selenium  with  the  members  of  the 
chlorine  group  present  no  features  of  special  interest  or 
importance.  In  general  they  resemble  the  correspond- 
ing compounds  of  sulphur,  but  are  more  stable.  No 
iodide  of  selenium  corresponding  to  the  hexiodide  of 
sulphur  has  been  discovered. 

TELLUKIUM,  (Te  At.  Wt.  126.52). 

Occurrence. — Tellurium  occurs  in  some  gold  ores  in 
the  native  or  uncombined  condition,  and  also  in  com- 
bination with  gold,  silver,  antimony,  lead,  and  other 
metals.  Tellurides  occur,  among  other  places  in  the 
United  States,  in  California  and  Virginia.  The  general 
method  of  preparing  tellurium  from  its  ores  is  the  same 
as  that  by  which  selenium  is  made.  The  tellurium  is 


HYDROGEN  TELLURIDE.  205 

oxidized  to  tellurious  acid,  H2TeO3,  which  is  isolated, 
and  then  reduced  by  means  of  sulphurous  acid. 

Properties. — Tellurium  is  silver-white,  and  crystallizes 
easily.  Heated  in  the  air  it  burns,  forming  a  thick 
white  cloud  of  tellurium  dioxide,  TeO2.  Treated  with 
sulphuric  acid  it  is  oxidized  to  tellurious  acid,  H2TeO3, 
the  sulphuric  acid  being  reduced.  Nitric  acid  oxidizes 
It  likewise  to  tellurious  acid.  Melted  together  with 
potassium  nitrate  it  is  converted  into  potassium  tellurate, 
K,TeO, 

Hydrogen  Telluride,  H2Te. — This  compound  is  made 
by  treating  zinc  telluride,  ZnTe,  with  an  acid.  It  is  a 
gas,  which  resembles  hydrogen  sulphide  in  most  of  its 
properties. 

Like  sulphur  and  selenium,  tellurium  combines  with 
the  members  of  the  chlorine  group.  The  compounds 
are  more  stable  than  those  of  the  other  two  elements  of 
the  group.  Thus,  taking  the  three  tetrachlorides,  that 
of  sulphur  is  extremely  unstable,  being  capable  of  exist- 
ence only  at  low  temperatures  ;  that  of  selenium  is  more 
stable,  but  still  it  is  decomposed  when  heated  to  the 
boiling  temperature ;  while  that  of  tellurium  can  be 
heated  far  above  the  temperature  of  boiling  without  de- 
composition. 

When  tellurium  tetrabromide  is  treated  with  potas- 
sium chloride  a  compound  of  the  formula  K2TeBr6  is 
formed,  which  in  composition,  as  will  be  observed,  is 
analogous  to  potassium  tellurite,  K2Te03,  two  atoms  of 
bromine  taking  the  place  of  each  atom  of  oxygen.  This 
appears  to  be  a  derivative  of  the  acid  H2TeBr6,  which  is 
similar  in  composition  to  the  chlorplatinic  acid,  H2PtCl6, 
referred  to  on  p.  142,  and  to  fluosilicic  acid,  H2SiF6.  As 
has  already  been  remarked,  not  many  such  compounds 
of  the  acid-forming  elements  are  commonly  met  with, 
largely  for  the  reason  that  they  are  for  the  most  part 
easily  decomposed  by  water  and  converted  into  the  cor- 
responding oxygen  compounds. 


CHAPTER.  XIV. 

COMPOUNDS  OF  SULPHUR,   SELENIUM,  AND  TELLURIUM 
WITH  OXYGEN  AND  WITH   OXYGEN  AND  HYDROGEN. 

Introductory. — It  has  already  been  stated  that,  when 
the  three  elements  of  the  sulphur  group  are  burned  in 
the  air,  they  are  converted  into  the  corresponding  diox- 
ides. Under  certain  conditions  it  is  possible  to  convert 
sulphur  dioxide  and  tellurium  dioxide  into  the  trioxides 
SO3  and  TeO3,  while  the  corresponding  compound  of 
selenium  has  not  been  made.  A  lower  oxide  of  sul- 
phur, S2O3,  and  a  higher  one  of  the  formula  S2O7  have 
also  been  made.  The  dioxides  dissolve  in  water,  and 
from  these  solutions  salts  of  the  general  formula  M2SO3r 
M2SeO3,  and  M2TeO3  are  obtained.  The  trioxide  of  sul- 
phur dissolves  in  water  with  great  evolution  of  heat, 
forming  compounds  S(OH)6,  OS(OH)4,  and  O2S(OH)2. 

Most  of  the  salts  obtained  from  this  solution  are  derived 
from  an  acid  of  the  formula  H2SO4,  and  therefore  this 
is  generally  called  sulphuric  acid.  By  treating  sulphuric 
acid  with  reducing  agents  of  various  kinds  it  can  be  con- 
verted successively  into  other  acids  containing  a  smaller 
proportion  of  oxygen ;  and  if  the  reduction  is  pushed 
far  enough  sulphur  and  hydrogen  sulphide  are  obtained. 
The  limit  of  reduction  is  reached  in  hydrogen  sulphide  ; 
and,  on  the  other  hand,  if  hydrogen  sulphide  is  oxidized 
the  limit  of  oxidation  is  reached  in  sulphuric  acid.  The 
intermediate  products  which  have  been  isolated  are 
represented  in  the  following  table,  by  the  side  of  which 
for  the  sake  of  comparison  is  placed  the  similar  series 
of  chlorine  acids  : 

H2S  HC1 

HC10 

H2S2O4  HC10, 

H2S03  HC108 

H,SO4  HC104 

(206) 


COMPOUNDS  OF  SULPHUR.  207 

In  the  chlorine  series,  however,  one  member  more  is 
known  than  in  the  sulphur  series.  A  remarkable  fact 
which  has  not  yet  been  explained  is  that  in  each  of  the 
ordinary  forms  of  the  acids,  both  in  the  chlorine  series 
and  in  the  sulphur  series,  the  number  of  hydrogen  atoms 
is  the  same  as  in  the  hydrogen  compound.  Throughout 
the  chlorine  series  there  is  but  one  hydrogen  atom  in 
the  molecule ;  throughout  the  sulphur  series  there  are 
two.  This  refers  only  to  those  forms  of  the  acid  from 
which  most  of  the  salts  are  derived,  and  plainly  not  to 
such  compounds  as  OS(OH)4  and  S(OH)6,  which,  as  we 
shall  see,  are  closely  related  to  sulphuric  acid,  just  as 
OC1(OH)5  or  H5C1O6  is  closely  related  to  perchloric  acid. 
The  number  of  compounds  of  sulphur  is  increased  by 
the  power  the  element  possesses  of  uniting  with  itself, 
much  as  it  does  with  oxygen.  Thus  the  sulphides  take 
up  sulphur  and  are  converted  into  polysulphides,  and 
the  limit  of  action  of  this  kind  is  a  compound  of  the  gen- 
eral formula  M2SS4.  This  series  is  complete,  the  mem- 
bers having  the  composition  represented  in  the  follow- 
ing table : 

M2S  M2S 

M2SS 

M2SS2 

M2SS3  M2S03 

M2SS4  M2SO4. 

Comparing  this  with  the  series  of  oxygen  compounds 
an  analogy  is  apparent,  though  the  second  and  third 
members  of  the  oxygen  series  are  lacking.  The  limit  is 
reached  in  a  compound  of  the  same  order  in  both  series. 
Again,  sulphur  can  take  the  place  of  part  of  the  oxy- 
gen in  the  oxygen  acids.  Thus  compounds  of  the  com- 
position represented  by  the  following  formulas  are  con- 
ceivable :  .'  -; 

H2SOS    or  H2S20 ; 

H2S02S  or  H2S202 ; 

H2SOS2  or  H2S3O ; 

H2S03S  or  H2S203 ; 

H2S02S2  or  H2S302 ; 

H2SOS3  or  H2S40. 


208  INORGANIC  CHEMISTRY. 

Of  these  compounds  only  that  one  which  has  the  compo- 
sition H2S2O3  is  known.  This  is  thiosulphuric  acid  or,  as 
it  is  sometimes  called,  hyposulphurous  acid. 

Other  complications  are  occasioned  by  the  combina- 
tion of  sulphur  with  sulphur  in  some  of  the  oxygen  acids 
above  mentioned.  In  this  way  probably  is  formed  the 
series  represented  in  the  following  table  : 

Dithionic    acid, H2S2O6 

Trithionic      "     .     .     ;    .   .V  ';•;  I'  H2S3O6 

Tetrathionic  " H2S4O6 

Pentathionic  "     .     ...     .     .     .'  '.  H2S5O6. 

While  the  number  of  compounds  which  sulphur  forms 
with  hydrogen  and  oxygen  is  comparatively  large,  the 
relations  between  them  can  be  traced  without  serious 
difficulty,  and,  studied  in  the  right  way,  they  are  seen 
to  follow  certain  laws  which  govern  their  composition. 
These  relations  will  be  discussed  after  the  main  facts 
concerning  the  principal  compounds  have  been  studied. 

The  number  of  compounds  which  selenium  and  tellu- 
rium form  with  hydrogen  and  oxygen  is  much  smaller 
than  in  the  case  of  sulphur.  The  only  ones  known  are 
those  which  are  analogous  in  composition  to  sulphurous 
acid,  H2SO3,  and  sulphuric  acid,  H2SO4 ;  besides  these, 
however,  there  are  a  few  compounds  known  which  con- 
tain hydrogen,  sulphur,  selenium,  and  oxygen.  These 
are  closely  related  to  some  of  the  compounds  of  sulphur 
with  hydrogen  and  oxygen,  being  derived  from  them  by 
the  introduction  of  selenium  for  a  part  of  the  oxygen. 

Sulphur  combines  with  oxygen  and  members  of  the 
chlorine  group,  and  also  with  these  elements  and  hydro- 
gen. The  simplest  compound  of  this  kind  is  that  which 
has  the  composition  SOC12.  This  is  plainly  analogous 
to  sulphur  dioxide,  from  which  it  can  be  made  by  sub- 
stituting two  atoms  of  chlorine  for  one  atom  of  oxygen  in 
the  molecule.  While  this  compound  cannot  unite  directly 
with  more  chlorine,  a  compound  of  the  composition 
SO2Cla  can  be  made.  This  is  in  accordance  with  what 
we  should  expect  from  a  consideration  of  the  conduct 
of  sulphur  towards  oxygen  and  towards  chlorine  respec- 


SULPHURIC  ACID.  209 

tively.  With  oxygen  it  forms  a  stable  compound,  SOS, 
while  the  limit  of  combination  with  chlorine  is  reached 
in  the  compound  SC14,  and  even  this  is  very  unstable. 
As  we  commonly  say,  sulphur  is  sexivalent  towards 
oxygen  and  quadrivalent  towards  chlorine.  Towards 
both  together  it  is  also  sexivalent,  as  shown  in  the  com- 
pound SO2C12,  though  the  compound  SOC14  does  not 
exist.  Just  as  the  simple  compounds  of  sulphur  and 
chlorine  are  decomposed  readily  by  water,  forming  oxy- 
gen compounds,  so  these  mixed  compounds  containing 
oxygen  and  chlorine  are  also  readily  decomposed  by 
water.  The  decomposition  appears  to  take  place  as 
represented  in  the  following  equations  : 


In  the  first  case,  the  product  formed  breaks  down  into 
sulphur  dioxide  and  water,  so  that  the  main  action  is 
represented  thus  : 


+  H20  =  S02  +  2HC1, 

and  the  change  therefore  consists  in  substituting  one 
atom  of  oxygen  for  the  two  atoms  of  chlorine.  It  is 
analogous  to  the  transformation  of  sulphur  tetrachloride 
into  sulphur  dioxide  : 

SC14  +  2H20  =  S02  +  4HC1. 

As  these  mixed  chlorides  and  oxides  are  readily  con- 
verted into  the  corresponding  acids  by  treatment  with 
water,  they  are  frequently  spoken  of  as  the  chlorides  of 
the  acids,  or  the  acid  chlorides.  Between  them  and  the 
anhydrides,  or  acidic  oxides,  there  is  plainly  an  analogy. 

Sulphuric  Acid,  H2SO4.  —  The  many  salts  of  sulphuric 
acid  can  be  best  explained  on  the  assumption  that  the 
forms  of  sulphuric  acid  are  derived  from  a  compound 
S(OH)6,  just  as  the  salts  of  periodic  acid  can  be  best  ex- 
plained on  the  assumption  of  a  fundamental  compound, 
I(OH)7.  As  the  latter  is  called  normal  periodic  acid,  so 


210  INORGANIC  CHEMISTRY. 

the  former  is  called  normal  sulphuric  acid.  From  normal 
sulphuric  acid  by  successive  losses  of  water  are  formed 
the  compounds  H4SO5,  H2SO4,  and  SO3 : 

S(OH)6  =  OS(OH)4  +  H2O  ; 
OS(OH)4  =  02S(OH)2  +  H20; 
02S(OH)2  =  03S  +  H20. 

While  these  compounds  are  all  known,  the  salts  of  sul- 
phuric acid  are  for  the  most  part  derived  from  the  acid 
containing  two  hydrogen  atoms,  viz.,  O2S(OH)2.  In  some 
salts  two  of  the  hydrogen  atoms  in  normal  sulphuric 
acid  are  replaced  by  metals,  the  others  remaining.  Salts 
of  the  general  formula  S(OH)4(OM)2  are  thus  formed. 
These  are  generally  represented  as  containing  two 
molecules  of  water  of  crystallization,  thus:  M2SO4+2H2O. 
To  decide  between  the  two  views  is  at  present  impos- 
sible. In  the  first  place,  the  question  as  to  the  nature 
of  water  of  crystallization  must  be  answered  before  it 
can  be  said  whether  there  is  any  conflict  between  the 
two  views.  As  far  as  the  assumption  of  normal  sul- 
phuric acid  is  concerned,  it  can  only  be  said  that  similar 
assumptions  in  the  case  of  iodine  and  in  that  of  phos- 
phorus seem  to  be  well  founded,  and,  although  it  would 
be  difficult  to  furnish  proof  positive  of  the  reality  of  the 
relations,  the  above  suggestion  is  by  far  the  simplest 
which  has  been  made,  and  it  is  of  great  assistance  in 
dealing  with  the  acid  derivatives  of  the  elements  of 
Families  VII,  VI,  V,  and  IV. 

For  the  science  as  well  as  for  the  art  of  chemistry 
sulphuric  acid  is  of  fundamental  importance.  It  is 
used  daily  in  every  chemical  laboratory  and  in  every 
chemical  factory,  and  in  some  of  the  most  important 
branches  of  chemical  industry  enormous  quantities  of 
it  are  used.  In  consequence  of  the  large  demand  for 
the  acid  the  process  used  in  preparing  it  has  been 
studied  with  great  care,  and  it  has  reached  a  high  state 
of  perfection.  As  it  furnishes  an  excellent  example  of 
the  applications  of  the  facts  of  science  to  the  building 
up  of  an  industry,  it  will  be  studied  with  some  degree  of 
fulness. 


SULPHURIC  ACID.  211 

Sulphuric  acid  has  been  known  for  a  long  time.  It 
was  made  in  the  eighteenth  century  by  heating  calcined 
iron  vitriol  (ferrous  sulphate,  FeSO4)  with  sand,  and 
was  hence  called  oil  of  vitriol,  a  name  which  still  sur- 
vives. It  was  also  prepared  by  treating  sulphur  with 
saltpeter  (potassium  nitrate,  KNO3).  The  acid  was  first 
prepared  on  the  large  scale  in  England  with  the  use  of 
comparatively  large  chambers  lined  with  lead.  It  was 
known  as  English  sulphuric  acid,  and  is  still  called  by 
this  name. 

Sulphuric  acid  occurs  in  nature  in  the  form  of  salts  or 
sulphates,  such  as  calcium  sulphate  or  gypsum,  barium 
sulphate  or  heavy  spar,  and  others.  Although  these 
salts  occur  in  large  quantity,  the  acid  is  not  obtained 
from  them,  as  there  is  no  economical  way  of  substituting 
hydrogen  for  the  metal.  The  preparation  of  the  acid  by 
the  substitution  of  hydrogen  for  the  calcium  in  calcium 
sulphate,  CaSO4,  suggests  itself,  but  this  cannot  be  easily 
effected  by  any  substance  available  in  quantity.  Hydro- 
chloric and  nitric  acids  are  made  from  their  salts  which 
occur  in  nature,  the  former,  as  we  have  seen,  from 
sodium  chloride,  NaCl,  the  latter  from  saltpeter  or  po- 
tassium nitrate,  KNO3,  by  treating  with  sulphuric  acid. 
There. is,  however,  no  acid  which  acts  upon  the  sulphates 
as  sulphuric  acid  acts  upon  chlorides,  nitrates,  and  many 
other  salts. 

The  manufacture  of  sulphuric  acid  is  based  upon  the 
two  fundamental  facts  that  (1)  when  sulphur  is  burned 
it  is  converted  into  sulphur  dioxide,  SO2 ;  and  (2)  when 
sulphur  dioxide  is  treated  with  an  oxidizing  agent  in  the 
presence  of  water  it  is  converted  into  sulphuric  acid : 

SO2      +  H2O  +  O  =  H2SO4 ;  or 
SO2      +  H2O  =  H2SO3 ;  and 

H2S03  +  0  =  H2S04. 

The  chief  difficulty  is,  of  course,  experienced  in  effect- 
ing the  oxidation  of  the  sulphurous  acid.  It  is  accom- 
plished by  introducing  the  gas,  sulphur  dioxide,  into 
large  chambers  together  with  compounds  of  nitrogen 


212  INORGANIC  CHEMISTRY. 

and  oxygen,  and  steam.  That  which  plays  the  chief 
part  in  effecting  the  transformation  is  a  mixture  of  the 
two  oxides  NO  and  NO2.  This  acts  like  the  trioxide,. 
N2O3,  and  it  may  be  represented  by  this  formula. 
Instead  of  starting  with  the  trioxide,  nitric  acid  is  used 
in  the  manufacture  of  sulphuric  acid.  This  at  first 
reacts  with  sulphur  dioxide  and  steam,  as  represented  in. 
the  equation 

2HNO3  +  2S02  +  H20  =  2H2SO4  +  N,O3- 

After  this  the  main  reactions  are  (1)  the  formation  of  a, 

NO 
compound  of  the  formula  SO2<Q-rr2,  called  nitrosyl-am- 

phuric  acid;  and  (2)  the  decomposition  of  the  nitrosyl- 
sulphuric  acid  by  water.  These  reactions  are  repre- 
sented in  the  two  equations  following  : 

280,  +  N.O,  +  O,  +  H,0  =  2SO,(OH)(NO,)  ; 


The  nitrogen  trioxide  formed  in  the  second  reaction 
then  again  enters  into  combination  with  sulphur  dioxide, 
oxygen,  and  water  to  form  nitrosyl-  sulphuric  aci$,  which 
again  undergoes  decomposition  with  water.  It  will  be 
seen,  therefore,  that  the  trioxide  is  not  lost,  but  that 
it  simply  serves  the  purpose  of  effecting  the  oxidation  of 
the  sulphur  dioxide,  and,  theoretically,  a  small  quantity 
of  the  gas  should  be  capable  of  transforming  an  infinite 
quantity  of  sulphur  dioxide  into  sulphuric  acid. 

Other  reactions  besides  those  mentione/i  above  are 
undoubtedly  involved  in  the  manufacture  of  sulphuric 
acid,  but,  according  to  the  most  elaborate  researches 
made  on  the  subject  in  a  factory  in  operation,  those 
mentioned  are  the  principal  ones.  Whatever  the  details. 
may  be,  the  oxidation  is  effected  without  difficulty,  and 
the  waste  in  nitrogen  compounds  is  now  but  slight. 

The  reactions,  as  has  been  said,  are  carried  on  in  large 
chambers  lined  with  lead,  and  known  as  the  leaden 
chambers.  The  sulphur  is  burned  in  a  special  furnace 


SULPHURIC  ACID. 


213 


so  arranged  that  air  has  free  access  to  it.  The  dioxide 
thus  formed  is  then  conducted  through  a  tower  so  con- 
structed that  it  presents  a  large  surface  to  the  action  of 
the  gas.  Through  this  tower  dilute  sulphuric  acid,  taken 
from  another  tower  at  the  end  of  the  system,  flows  from 


above,  and  the  hot  gases  coming  in  contact  with  this 
serve  the  purpose  of  concentrating  it,  while  the  gas  itself 
is  cooled  before  entering  the  leaden  chamber.  The 
general  arrangement  of  the  essential  parts  of  a  sulphuric 
acid  factory  are  shown  in  Fig.  7. 


214  INORGANIC  CHEMISTRY. 

The  sulphur  dioxide  passes  through  the  large  tube  x 
into  the  tower  G,  called  the  Glover  tower.  This  is  filled 
from  d  to  e  with  pieces  of  fire-brick  over  which  from  the 
cistern  b  a  continual  stream  of  dilute  sulphuric  acid 
flows.  The  gases  are  thus  cooled  down  and  the  acid 
concentrated.  From  a  second  cistern  there  flows  at  the 
same  time  concentrated  sulphuric  acid  from  the  tower 
G',  or  the  Gay  Lussac  tower.  This  stronger  acid  con- 
tains oxides  of  nitrogen  in  combination,  and  by  contact 
with  the  dilute  acid  the  oxides  are  set  free  and  are  thus 
mixed  with  the  sulphur  dioxide.  The  nitric  acid  is  in- 
troduced into  the  first  chamber,  No.  1,  in  the  form  of 
vapor,  together  with  the  oxides  of  nitrogen  and  sulphur 
dioxide,  and  here  also  the  gases  meet  with  water  vapor. 
The  reactions  above  referred  to  now  take  place.  From 
the  first  chamber  the  gases  pass  through  the  pipe  v  into 
the  second,  from  this  into  the  third  at  w,  and  finally, 
in  order  to  prevent  the  escape  of  any  unused  oxides 
of  nitrogen,  the  gases  are  passed  through  the  Gay 
Lussac  tower  G'.  This  contains  coke  over  which  is 
kept  flowing  concentrated  sulphuric  acid  which  takes  up 
the  oxides  of  nitrogen  and  will  give  them  up  again  when 
diluted.  This  liberation  of  the  oxides  of  nitrogen  is 
accomplished,  as  has  been  said,  in  the  Glover  tower. 
The  concentrated  acid  collected  at  the  bottom  of  the 
Glover  tower  is  well  adapted  for  use  in  the  Gay  Lussac 
tower.  The  leaden  chambers  in  some  factories  are 
nearly  100  feet  long,  18  to  30  feet  broad,  and  15  feet 
high. 

The  acid  taken  from  the  chambers  contains  about  64 
per  cent  of  the  compound  H2SO4,  and  has  the  specific 
gravity  1.5.  This  is  evaporated  first  in  lead  pans  until 
it  reaches  the  specific  gravity  1.75.  As  stronger  acid 
acts  upon  lead,  the  evaporation  beyond  this  point  is 
carried  on  in  platinum,  glass,  or  iron.  The  strong  acid 
thus  obtained,  which  has  a  specific  gravity  of  about  1.830, 
is  the  concentrated  sulphuric  acid  of  commerce. 

Instead  of  sulphur,  iron  pyrites  is  now  extensively 
used  for  the  preparation  of  sulphur  dioxide  in  the  manu- 
facture of  sulphuric  acid.  This  is  a  compound  of  iron 


SULPHURIC  ACID.  215 

and  sulphur  of  the  composition  FeS2.  When  it  is  heated 
in  contact  with  the  air  it  is  converted  into  the  oxide,  and 
the  sulphur  passes  off  in  the  form  of  sulphur  dioxide. 
If  the  sulphide  used  contains  selenides  the  selenium 
dioxide  formed  in  the  roasting  process  is  carried  into 
the  flues,  and  is  there  deposited  with  the  flue-dust,  as 
was  stated  in  speaking  of  the  source  of  selenium. 

Commercial  sulphuric  acid  is  an  oily  liquid,  usually 
somewhat  colored  by  impurities.  The  nature  of  the 
impurities  is  dependent  upon  the  substances  used  in  the 
manufacture  and  upon  the  conditions.  It  often  con- 
tains some  lead  sulphate  in  solution,  and  when  it  is 
diluted  with  water  this  separates,  giving  the  liquid  a 
more  or  less  cloudy  appearance.  By  standing,  however, 
the  liquid  becomes  clear,  as  the  lead  sulphate  collects 
at  the  bottom.  For  obvious  reasons,  some  oxides  of 
nitrogen  are  also  generally  present.  Among  other  im- 
purities frequently  met  with  in  the  commercial  acid  are 
arsenic  from  the  pyrites,  and  a  little  selenium. 

Pure  Sulphuric  Acid  is  made  from  the  commercial  acid 
by  treating  it  with  such  substances  as  will  remove  the 
oxides  of  nitrogen  and  arsenic,  and  by  distilling.  It  is 
a  colorless  liquid  at  the  ordinary  temperature.  The 
pure  acid  generally  made  has  about  the  same  concentra- 
tion as  the  commercial  crude  acid.  By  taking  special  pre- 
cautions in  the  distillation  a  product  having  very  nearly 
the  composition  H2SO4  can  be  obtained.  This  is  a  thick, 
clear  liquid  of  specific  gravity  1.854  at  0°.  When  cooled 
down  to  a  low  temperature  it  forms  large  crystals  which 
melt  at  10°. 5.  When  heated  it  gives  off  some  sulphur 
trioxide  in  consequence  of  partial  decomposition  into 
this  substance  and  water  : 

H2S04  =  H20  +  S03. 

It  finally  boils,  however,  at  the  temperature  338°,  and 
the  distillate  has  the  composition  represented  by  the 
formula  H2SO4  +  -^-HaO.  Heated  somewhat  above  its 
boiling  point  it  is  completely  decomposed  into  sulphur 
trioxide  and  water,  according  to  the  above  equation.  If 
the  heating  is  carried  to  a  higher  temperature  the  sul- 


216  INORGANIC  CHEMISTRY. 

phur  trioxide  breaks  down  into  sulphur  dioxide  and 
oxygen. 

Sulphuric  acid  has  a  strong  tendency  to  absorb  water, 
and  to  form  compounds  with  it.  In  consequence  of  the 
formation  of  these  compounds  a  great  deal  of  heat  is 
evolved  when  sulphuric  acid  is  mixed  with  water.  One 
molecule  of  sulphuric  acid  when  mixed  with  about  1600 
molecules  of  water  gives  17,850  cal.  Among  the  com- 
pounds thus  formed  are  the  so-called  hydrates  of  the 
composition  H2SO4  +  H2O,  and  H2SO4  +  2H2O,  which 
should  probably  be  regarded  as  having  the  constitution 
OS(OH)4  and  S(OH)6. 

So  strong  is  the  tendency  of  the  acid  to  combine  with 
the  elements  of  water  that  it  abstracts  them  in  the  pro- 
portions to  form  water  from  many  organic  substances. 
Thus  a  piece  of  wood  placed  in  sulphuric  acid  soon 
turns  black  and  is  completely  disintegrated.  The  cause 
of  this  is  to  be  found  in  the  fact  that  wood  consists 
essentially  of  carbon,  hydrogen,  and  oxygen ;  and  the 
sulphuric  acid  abstracts  the  hydrogen  and  oxygen, 
leaving  the  carbon  mainly  in  the  uncombined  state. 
Similarly  it  abstracts  moisture  from  gases,  and  it  is  used 
extensively  in  the  laboratory  for  this  purpose.  In  con- 
tact with  the  skin  it  acts  as  upon  wood,  causing  wounds 
which  are  painful  and  difficult  to  heal. 

Sulphuric  acid  is  what  is  called  a  strong  acid,  a  term 
which  needs  further  explanation,  and  the  subject  of  the 
relative  strengths  of  acids  will  be  discussed  in  due  time. 
As  used  here,  it  means  simply  that  the  acid  has  the 
power  to  decompose  the  salts  of  most  other  acids,  ap- 
propriating the  metals  and  setting  the  other  acids  free. 
This  is  illustrated  by  the  formation  of  hydrochloric  and 
nitric  acids  by  treating  sodium  chloride  and  potassium 
nitrate  respectively  with  sulphuric  acid.  We  shall  see, 
however,  that  the  fact  that  sulphuric  acid  decomposes 
the  salts  of  hydrochloric  and  nitric  acids  does  not  prove 
that  it  is  a  stronger  acid  than  they  are.  There  are  other 
facts  besides  the  strengths  of  the  acids  which  determine 
whether  such  decompositions  take  place  or  not. 

Sulphuric  acid  gives  up  its  oxygen  to  other  substances 


SULPHURIC  ACID.  217 

comparatively  easily,  and  is  generally  reduced  to  sul- 
phurous acid  which  is  decomposed  into  sulphur  dioxide 
and  water.  Thus  hydrogen  sulphide  and  hydrobromic 
acid  both  act  upon  it,  as  we  have  seen ;  the  products  be- 
ing, in  the  former  case,  sulphur  dioxide,  water,  and  sul- 
phur ;  in  the  latter,  sulphur  dioxide,  water,  and  bromine : 

H2S04  +  H2S     =  2H20  +  SO2  +  S ; 
H2S04  +  2HBr  =  2E2O  +  SO2  +  Br2. 

Generally,  in  acting  upon  metals  it  forms  the  correspond- 
ing salt  and  hydrogen  is  given  off,  but  if  the  temperature 
is  high  and  the  acid  concentrated  the  hydrogen  acts 
upon  the  acid,  reducing  it.  Thus,  in  making  hydrogen 
by  treating  zinc  with  sulphuric  acid,  if  the  acid  is  dilute 
and  is  kept  cool  the  reaction  takes  place  in  the  simplest 
way  ;  but  if  the  acid  is  concentrated  and  is  allowed  to 
get  hot  some  hydrogen  sulphide  is  always  formed. 
Copper  does  not  act  upon  sulphuric  acid  at  ordinary 
temperatures.  If  it  is  treated  with  concentrated  sul- 
phuric acid,  sulphur  dioxide  is  the  chief  reduction- 
product.  Even  free  hydrogen  if  passed  into  sulphuric 
acid  heated  to  160°  reduces  it,  forming  sulphur  dioxide. 
The  complete  reduction  of  sulphuric  acid  to  sulphur  and 
hydrogen  sulphide  is  beautifully  shown  by  the  action  of 
hydriodic  acid.  As  was  stated  in  speaking  of  the  action 
of  sulphuric  acid  upon  potassium  iodide,  four  reactions 
may  take  place  when  these  substances  act  upon  each 
other.  They  are  represented  by  the  equations 

2KI     +  H2S04  =  K2SO4  +  2HI ; 
H2S04  +  2HI     =  2H20  +  S02  +  21 ; 
H2S04  +  6HI     =4H20  +  S     +61; 
H2SO4  +  8HI     =4H20  +H.S  +  8I. 

Carbon  and  sulphur  also  act  upon  sulphuric  acid  and 
reduce  it  to  sulphur  dioxide.  With  sulphur  the  action 
is  represented  thus : 

2H2S04  +  S  =  3S02  +  2HSO. 


218  INORGANIC  CHEMISTRY. 

Sulphuric  acid  is  a  dibasic  acid  (see  p.  138)  and  forms 
two  series  of  sajts,  normal  salts  of  the  general  formula 
M^O^  and  acid  salts  of  the  formula  MHSO4.  The  sul- 
phates are  very  stable  salts.  Those  of  the  strongest 
bases  are  not  decomposed  by  the  highest  heat.  Those  of 
weaker  bases  break  down,  giving  up  sulphur  trioxide  ;  or 
if  the  decomposition  takes  place  at  a  temperature  above 
that  at  which  sulphur  trioxide  breaks  down,  this  de- 
composes into  sulphur  dioxide  and  oxygen. 

Tetrah.ydroxyl-Sulph.uric  Acid,  OS(OH)4.  —  This  com- 
pound has  already  been  referred  to  as  being  formed  by 
the  addition  of  water  to  ordinary  sulphuric  acid.  It  is 
a  crystallized  compound  which  melts  at  7.5°.  No  salts 
of  this  acid  are  known,  or,  rather,  no  salts  derived  from 
it  by  the  replacement  of  all  the  hydrogen  by  metal  are 
known.  Only  two  of  the  hydrogen  atoms  appear  to  be 
replaceable.  This  compound  is  generally  represented 
by  the  formula  H2SO4  -f-  H2O,  and  called  a  hydrate. 

Normal  Sulphuric  Acid,  S(OH)8.  —  This  is  commonly 
represented  by  the  formula  H2SO4  -j-  2H2O  and,  like  the 
preceding  compound,  regarded  as  a  hydrate.  It  appears 
to  be  formed  by  the  action  of  water  on  sulphuric  acid. 
On  mixing  sulphuric  acid  and  water  the  maximum  con- 
traction takes  place  when  the  quantities  necessary  to 
form  this  compound  are  brought  together,  and  there  are 
other  good  reasons  for  believing  that  the  compound 
exists  in  the  solution.  It  is  not  a  solid  like  the  pre- 
ceding compound.  The  acid  forms  no  salts  bearing  sim- 
ple relations  to  it.  There  are  some  salts  known,  how- 
ever, which  appear  to  be  related  to  it,  as,  for  example, 
the  salt  K3HS2O8  and  some  others  similar  to  it.  The 
acid  from  which  these  salts  are  derived  is  believed  to  be 
formed  from  the  normal  acid  by  elimination  of  water,  as 
represented  below  : 

fOH      HOI 

18>  S1  OH  +  HO*  S<OH  =  (HO)2OS<g>SO(OH)2  +  4H20 
OH       HO  I 


An   acid  formed  in   this  way  would  have  the  formula 


DISULPHURIC  ACID.  219 

H4S2Og,  and  the  salt  K3HS2O8  is  the  tertiary  potassium 
salt  of  this  acid. 

Disulphuric  Acid,  Pyrosulphuric  Acid,  H2S2O7. — This 
compound,  which  is  also  known  by  the  names  fuming 
sulphuric  acid  and  Nordhausen  sulphuric  acid,  is  closely 
related  to  ordinary  sulphuric  acid,  and  is  made  from  it 
by  treating  it  with  sulphur  trioxide,  the  two  combining 
directly,  as  represented  thus  : 

H2SO4  +  SO3  =  H2SaO7. 

It  is  made  by  distilling  ferrous  sulphate  which  is  not 
perfectly  dry  : 

4FeS04  +  H20  =  2FeaO3  +  2SO2  +  HaSa07. 

A  so-called  solid  sulphuric  acid  is  now  manufactured  by 
a  process  which  will  be  referred  to  under  Sulphur  Triox- 
ide. It  is  essentially  disulphuric  acid. 

Disulphuric  acid,  as  it  is  found  in  the  market,  is  gen- 
erally a  thick  liquid  which  gives  off  dense  fumes  when 
exposed  to  the  air,  and  breaks  down  completely  into  sul- 
phur trioxide  and  sulphuric  acid  when  heated.  When 
pure  it  crystallizes  in  large  crystals  which  melt  at  35°. 

It  is  believed  that  the  relation  between  disulphuric 
acid  and  ordinary  sulphuric  acid  should  be  expressed  by 
these  formulas : 


*ro>so  =  0  s-  -, 

H0  \OH  HO/ 


Or  the  formula  of  the  acid  may  be  written  thus  : 

02S-OH 


This  relation  is  similar  to  that  believed  to  exist  between 
the  normal  acid  S(OH)6  and  the  acid  H4S2O8  (see  p.  218). 


220  INORGANIC  CHEMISTRY. 

Disulphuric  acid  forms  normal  salts  of  the  general  for- 
mula M2S2O7  and  acid  salts  of  the  general  formula 
MHS2O7.  When  heated  the  normal  salts  break  down, 
yielding  ordinary  sulphates  and  sulphur  trioxide  : 

M2S2O7  =  M2SO4  +  S03. 

When  an  acid  sulphate  like  KHSO4  is  heated  to  a 
sufficiently  high  temperature  it  breaks  down  into  water 
and  a  disulphate  : 

2KHS04  =  K2S207  +  H20. 

Sulphurous  Acid,  H2SO3. — While  no  acid  of  the  formula 
H2SO3  is  known  in  the  free  condition,  a  large  number  of 
salts  which  are  derived  from  this  acid  are  known.  They 
are  made  by  treating  a  water  solution  of  sulphur  dioxide 
with  bases,  and  therefore  it  is  believed  that  the  solution 
contains  the  acid  which  is  formed  by  the  action  of  sulphur 
dioxide  on  water,  thus : 

S02  +  H20  =  H2S03. 

It  is,  however,  so  unstable  that  it  breaks  down  into  the 
dioxide  and  water  at  every  attempt  to  isolate  it.  The 
dioxide,  as  has  been  stated  and  as  will  be  shown  more 
fully,  is  formed  by  the  burning  of  sulphur  and  by  the 
reduction  of  sulphuric  acid.  The  acid  forms  a  number 
of  unstable  hydrates,  apparently  of  complicated  com- 
position. Owing  to  their  great  instability,  however,  the 
Investigation  of  these  substances  is  extremely  difficult 
and  unsatisfactory. 

Sulphurous  acid  takes  up  oxygen  readily  and  is  thus 
transformed  into  sulphuric  acid.  It  is  only  necessary  to 
allow  a  solution  to  stand  for  a  time  to  find  that  the  odor 
of  the  gas  disappears  and  that  sulphuric  acid  is  then 
present  in  the  solution.  Sulphurous  acid  is  frequently 
used  in  the  laboratory  as  a  reducing  agent.  Its  action  in 
this  way  has  been  illustrated  in  the  method  for  the  ex- 
traction of  selenium  from  selenious  acid  (see  p.  203).  The 
reaction  in  this  case  is  represented  thus  : 


SULPHUROUS  ACID.  221 

H2Se03  +  2S02  +  H2O  =  Se  +  2H2SO4 ;  or 
H2SeO3  +  2H2S03  =  Se  +  2H2SO4  +  H2O. 

Another  illustration  of  its  power  as  a  reducing  agent  is 
-shown  in  its  action  upon  iodic  acid.  When  it  is  added 
to  a  solution  containing  iodic  acid,  HIO3,  iodine  separates, 
the  reaction  taking  place  in  accordance  with  the  follow- 
ing equation : 

2HI03  +  5H2SO8  =  H20  +  5H2SO4  +  I2. 

If  sulphurous  acid  is  added  to  the  liquid  in  which  the 
iodine  is  suspended  further  action  takes  place,  the  iodine 
being  reduced  to  hydriodic  acid,  thus  : 

H2S03  +  H20  •+  I2  =  H2S04  +  2HI. 

This  reaction  takes  place  only  in  dilute  solution.  Con- 
centrated sulphuric  acid  acts  upon  hydriodic  acid  and  is 
reduced  by  it,  as  we  have  seen.  Towards  some  sub- 
stances sulphurous  acid  acts  as  an  oxidizing  agent,  and 
is  itself  reduced  to  lower  forms,  as  hyposulphurous  acid, 
H2S,,O4,  and  hydrogen  sulphide.  This  has  been  illustrated 
in  the  action  of  hydriodic  acid  upon  sulphuric  acid.  It 
is  also  illustrated  in  the  action  of  zinc  upon  sulphurous 
acid  in  the  presence  of  hydrochloric  acid,  when  this  re- 
action takes  place  : 

3Zn  +  6HC1  +  H2SO3  =  3ZnCl2  +  3H2O  +  H2S. 

Treated  with  zinc  alone  the  acid  sodium  salt,  NaHSO3, 
is  reduced  as  shown  in  the  following  equation,  a  salt  of 
hyposulphurous  acid,  H2S2O4 ,  being  formed  : 

Zn  +  4NaHSO3  =  ZnSO3  +  Na2SO3  +  Na2S2O4  +  2H2O. 

Sulphurous  acid  when  heated  in  a  sealed  tube  breaks 
down  into  sulphuric  acid  and  sulphur  : 

3H2S03  =  2H2S04  +  H20  +  S. 
This  kind  of  decomposition  is  also  characteristic  of  the 


"222  INORGANIC  CHEMISTRY. 

salts  of  the  acid  with  the  alkali  metals,  as,  for  example,, 
sodium  sulphite  : 

4NaaSO3  =  3Na2S04  +  Na2S. 

In  fact,  whenever  an  alkali  salt  of  any  of  the  oxygen 
acids  of  sulphur  is  heated  the  tendency  is  towards  the 
formation  of  the  sulphate,  all  the  oxygen  in  the  salt 
going  to  form  sulphate  ;  and  the  other  elements  in  excess 
of  what  may  be  needed  for  the  sulphate  arranging  them- 
selves in  other  forms. 

Just  as  the  sulphites  take  up  oxygen  to  form  sul- 
phates, they  also  take  up  sulphur  to  form  thiosulphates. 
The  two  reactions  appear  to  be  perfectly  analogous  : 

Na2SO3  +  0=Na2S04; 

Na2SO3  +  S  =  Na2S2O3  (or  Na2SO3S). 

t 

Sulphurous  acid  forms  two  series  of  salts,  the  normal 
sulphites  of  the  general  formula  M2SO3,  and  the  acid 
sulphites  of  the  general  formula  MHSO3.  These  are  un- 
stable, though  not  as  markedly  so  as  the  acid  itself. 
When  treated  with  most  acids  they  are  decomposed,, 
yielding  sulphur  dioxide  instead  of  sulphurous  acid. 
The  decomposition  of  sodium  sulphite  with  hydrochloric 
acid  is  represented  by  the  equation 

Na2SO3  +  2HC1  =  2NaCl  +  H2O  +  SO, ; 
with  sulphuric  acid  thus  : 

Na2S03  +  H2S04  =  Na2SO4  +  H2O  +  SO2. 

The  sulphites,  like  sulphurous  acid,  combine  readily 
with  oxygen,  tending  to  pass  over  into  the  sulphates,  and, 
as  has  been  remarked,  they  also  tend  to  unite  with  sul- 
phur to  form  the  thiosulphates. 

Hyposulphurtfus  Acid,  H2S2O4. — This  acid  is  also  called 
hydrosulphurous  acid,  but  the  name  hyposulphurous 
acid  is  more  in  accordance  with  the  nomenclature  adopted 
for  the  other  acids,  and  is  now  preferred.  But  little  is 


THIOSULPHURIC  ACID.  223 

known  of  the  compound.  It  is  formed  by  reduction  of 
.a  salt  of  sulphurous  acid  by  treating  with  zinc,  when  this 
reaction  takes  place : 

Zn  +  4NaHS08  =  ZnSO8  +  NaaSO.  +  NaaSaO,  +  2HaO. 

The  free  acid  cannot  be  obtained  from  this  sodium  salt, 
.and  notwithstanding  the  fact  that  it  has  been  the  subject 
of  a  number  of  elaborate  investigations  there  is  still 
some  doubt  as  to  the  composition  of  the  salt,  the  dis- 
coverer still  claiming  that  it  has  the  composition  repre- 
sented by  the  formula  NaaSOa ,  according  to  which  it  is  to 
be  referred  to  an  acid  of  the  formula  H3SOa. 

Hyposulphurous  acid,  like  sulphurous  acid,  combines 
readily  with  oxygen,  and  passes  first  into  sulphurous  and 
then  into  sulphuric  acid.  Its  reducing  action  is  stronger 
than  that  of  sulphurous  acid.  It  is  decomposed  by 
standing  in  the  air,  yielding  first  thiosulphuric  acid, 
H2S2O3 ,  and  then  sulphur  dioxide,  sulphur,  and  water. 

Hyposulphurous  acid,  so  far  as  composition  is  con- 
cerned, occupies  a  position  between  thiosulphuric  acid, 
H2S2O3 ,  and  pyrosulphurous  acid,  H2SaO6 ,  which  latter  is 
related  to  sulphurous  acid  in  the  same  way  that  disul- 
phuric  or  pyrosulphuric  acid,  H2S2O7 ,  is  related  to  sul- 
phuric acid.  The  relations  between  hyposulphurous, 
thiosulphuric,  and  pyrosulphurous  acids  are  shown  in 
the  table  below  : 

H,S203; 
IW)4; 
HaS20, 

Thiosulphuric  Acid,  H2S2O3. — This  acid  was  formerly, 
and  is  still  by  many,  called  hyposulphurous  acid.  Its 
formation,  or  the  formation  of  its  salts  by  the  addition 
of  sulphur  to  the  sulphites,  has  been  mentioned,  and  the 
analogy  between  this  reaction  and  that  of  the  formation 
of  sulphates  by  the  addition  of  oxygen  to  sulphites  has 
been  commented  upon.  It  may  be  regarded  as  sulphuric 
acid  in  which  sulphur  has  been  substituted  for  one  atom 


224  INORGANIC  CHEMISTRY. 

of  oxygen,  and  hence  the  name  thiosulphuric  acid  is 
appropriate,  whereas  the  name  hyposulphurous  acid 
suggests  at  once  a  compound  similar  to  sulphurous  acid, 
but  containing  less  oxygen,  and  is  therefore  inappro- 
priate. Sodium  thiosulphate  is  formed  together  with  the 
pentasulphide  by  the  action  of  sulphur  upon  sodium 
hydroxide  : 

GNaOH  +  12S  =  2Na3S6  +  Na,SaO3  +  3HaO. 

The  sulphides  of  the  alkali  metals  pass  over  into  the 
corresponding  thiosulphates  by  the  action  of  oxygen. 
Thus  the  pentasulphide  is  changed  when  exposed  to  the 
air  in  aqueous  solution.  The  action  consists  in  a  sub- 
stitution of  three  atoms  of  oxygen  for  three  of  sulphur  : 

Na2S5  +  30  =  Na2S203  +  38. 

Sodium  thiosulphate  is  formed,  further,  by  the  action  of 
iodine  upon  a  mixture  of  sodium  sulphide  and  sodium 
sulphite  : 


21  =         >  2NaI. 


The  acid  itself  is  very  unstable,  breaking  down  into  sul- 
phur dioxide,  sulphur,  and  water  (see  p.  190).  By  acids  its 
salts  are  decomposed  in  a  similar  way  with  evolution  of 
sulphur  dioxide  and  separation  of  sulphur,  which  appears 
suspended  in  the  liquid  in  a  very  fine  state  of  division. 
With  hydrochloric  acid  the  decomposition  takes  place 
thus  : 

Na2S2O3  +  2HC1  =  2NaCl  +  SO2  +  S  +  H2O. 

When  heated  the  thiosulphate  of  sodium  breaks  down 
according  to  the  rule  stated  in  speaking  of  the  decompo- 
sition of  the  sulphite  by  heat.  All  the  oxygen  goes  to 
the  formation  of  the  sulphate,  and  the  elements  left  over 


OTHER  ACIDS  OF  SULPHUR.  225 

in  excess  of  what  is  required  for  the  sulphate  unite  to 
form  another  compound,  thus  : 

4Na2S2O3  =  3Na2SO4  +  Na2S5. 

Other  Acids  of  Sulphur.  —  Of  the  other  acids  of  sulphur 
but  little  need  be  said  here.  As  was  stated  on  page  208, 
these  acids  form  a  series  the  members  of  which  are  closely 
related  to  one  another.  The  series  is  as  follows  : 

Dithionic  acid,     .......  H2S2O6 

Trithionic  acid,    .......  H2S3O6 

Tetrathionic  acid,    ......  H2S4O6 

Pentathionic  acid,    .     .     .     .     .     .  H2S5O6 

Dithionic  Acid,  or  a  salt  of  the  acid,  is  made  by  passing 
sulphur  dioxide  into  water  having  finely  powdered  man- 
ganese dioxide  in  suspension.  This  reaction  takes  place  : 

MnO2  +  2SO2  =  MnS2O6. 

The  product  is  the  manganese  salt  of  dithionic  acid, 
and  from  this  other  salts  can  be  prepared.  The  free 
acid  breaks  down  readily  into  sulphuric  acid  and  sulphur 
dioxide  : 


So,  too,  when  a  salt  of  the  acid  is  heated  it  breaks  down, 
forming  a  sulphate  and  sulphur  dioxide  : 

K2S206  =  K2S04  +  S02. 

Trithionic  Acid,  H2S3Ofl,  or  its  potassium  salt,  is  formed 
by  treating  a  solution  of  acid  potassium  sulphite,  KHSO3, 
with  "flowers  of  sulphur,"  when  reaction  takes  place 
thus: 

6KHSO3  +  28  =  2K2S306  +  K2S2O3  +  3H2O. 
"When   the    dry    potassium    salt   is   heated   it   is   de- 


226  INORGANIC  CHEMISTRY. 

composed,  yielding  the  sulphate,  sulphur  dioxide,  and 
sulphur  : 


Similarly,  when  treated  with  acids,  decomposition  takes 
place  with  evolution  of  sulphur  dioxide,  separation  of 
sulphur,  and  formation  of  sulphuric  acid  : 

H2S306  =  H2S04  +  SO,  +  S. 

Tetrathionic  Acid,  HaS4O6,  is  made  from  salts  of  thio- 
sulphuric  acid  by  treating  them  with  iodine.  Thus  with 
sodium  thiosulphate  the  reaction  is 

2Na2S2O3  +  21  =  Na2S4O6  +  2NaI. 

The  acid  is  moderately  stable,  so  that  a  dilute  solution 
can  be  boiled  without  undergoing  decomposition.  When 
the  concentrated  acid  is  heated,  however,  it  breaks  down 
into  sulphuric  acid,  sulphur  dioxide,  and  sulphur  : 

H2S406  =  H2S04  +  S02  +  28. 

Pentathionic  Acid,  H2SBO6,  is  formed  by  the  action  of 
hydrogen  sulphide  upon  a  solution  of  sulphur  dioxide  in 
water. 

Persulphuric  Acid,  H2S3Oe,  is  formed  by  dissolving  the 
oxide  S2OT  in  water,  the  oxide  itself  being  formed  by 
subjecting  a  mixture  of  sulphur  dioxide  and  oxygen  to 
the  silent  discharge  in  an  ozone  tube  (see  p.  85).  The 
potassium  salt  of  persulphuric  acid  is  easily  obtained  by 
subjecting  to  electrolysis  a  saturated  solution  of  acid 
potassium  sulphate. 

Constitution  of  the  Acids  of  Sulphur.  —  The  existence 
of  the  oxide  SO,  and  of  the  iodide  SI,  seems  to  show  that 
sulphur  is  sexivalent  towards  oxygen  and  towards  iodine. 
Considering,  further,  the  facts  presented  under  the  head 
of  Periodic  Acid  (which  see),  which  can  only  be  explained 
satisfactorily  by  the  aid  of  the  assumption  that  the  differ- 
ent varieties  of  periodic  acid  are  derived  from  the  normal 
acid  I(OH),,  and  the  analogous  facts  presented  under  Sul- 


CONSTITUTION  OF  THE  ACIDS  OF  SULPHUR. 

phuric  Acid,  which  lead  to  the  belief  that  this  acid  is  de- 
rived from  the  normal  acid  S(OH)6,  the  arguments  in  favor 
of  the  sexivalence  of  sulphur  in  sulphuric  acid  are  seen 
to  be  strong,  though  not  conclusive.  On  the  other  hand,  if 
sulphur  is  sexivalent  in  sulphur  trioxide  and  in  sulphuric 
acid,  it  is  quadrivalent  in  sulphur  dioxide  and  sulphur 
tetrachloride,  and  bivalent  in  hydrogen  sulphide  and  sul- 
phur dichloride.  But  if  sulphur  is  sexivalent  in  sul- 
phuric acid  the  constitution  of  the  acid  must  be  repre- 

OH 


rented   by   the   formula  H-O-S-O-H  or   S4   °H.     Of 


O 

course  such  a  formula  as  this  involves  the  hypothesis 
that  when  an  oxygen  atom  is  combined  with  only  one 
other  atom  two  valences  or  affinities  are  brought  into 
play,  and  in  regard  to  this  we  have  very  little  if  any  evi- 
dence. It  may  be  said,  however,  that  if  oxygen  which 
is  thus  combined  is  replaced  by  univalent  atoms  its 
place  is  always  taken  by  two  of  these,  indicating  that 
whatever  the  power  may  be  which  holds  the  oxygen  atom 
in  combination,  that  same  power  can  hold  two  chlorine 
atoms,  etc.,  and  it  is  convenient  to  use  the  double  line  to 
indicate  the  existence  of  this  power. 

The  view  expressed  by  the  above  formula  in  regard  to 
the  structure  of  sulphuric  acid  has  been  tested  experi- 
mentally by  methods  which  appear  somewhat  compli- 
cated, but  the  principle  involved  can  be  easily  explained. 
If  the  formula  is  correct,  then  both  hydroxyl  groups  bear 
the  same  relation  to  the  sulphur,  and  so  also  do  the  two 
hydrogen  atoms.  Whether  one  or  the  other  of  these  hy- 
drogen atoms  be  replaced  by  another  atom  or  group  of 
atoms,  the  same  product  should  result.  Or,  if  one  of  the 
hydrogen  atoms  is  replaced  by  one  group  and  the  other 
by  another  group,  it  should  make  no  difference  in  which 
order  the  two  groups  are  introduced.  The  product  should 
be  the  same  in  the  two  cases.  Thus,  suppose  one  hydro- 
gen atom  is  replaced  by  a  group  X,  and  the  other  by  Y, 
the  product  should  be  represented  by  the  formula 


228  INORGANIC  CHEMISTRY. 

o  o 

II  II 

Y-O-S-O-X  in  one  case,  and  by  X-O-S-O-1   in  the 

II  II 

O  O 

other  case.  But  if  the  formula  given  for  sulphuric  acid 
is  correct  the  two  compounds  are  identical.  By  methods 
which  involve  the  use  of  apparently  complex  organic 
compounds  the  two  hydrogen  atoms  have  been  thus  re- 
placed by  two  different  groups,  first  in  one  way  and  then 
in  the  reverse  order ;  and  the  two  products  have  been 
found  to  be  identical.  Further,  it  has  been  shown  that 
when  the  two  hydroxyl  groups  of  sulphuric  acid  are  re- 

:C1 

placed  by  chlorine,  forming  the  compound  SK  Cl,  and 

the  chlorine  atoms  then  replaced  by  certain  groups  of 
atoms,  these  groups  are  in  direct  combination  with  sul- 
phur and  not  with  oxygen.  All  the  evidence  points  to 
the  conclusion  that  the  view  represented  above  is  correct, 
In  attempting  to  determine  the  constitution  of  sulphur- 
ous acid  a  new  difficulty  arises.  Just  as  hydrogen  which 
is  in  combination  with  oxygen  is  replaceable  by  metals, 
or  is  acidic,  so,  also,  is  hydrogen  which  is  in  combination 
with  sulphur.  It  is  possible,  therefore,  to  conceive  of 
two  arrangements  of  the  atoms  composing  sulphurous 
acid,  both  representing  dibasic  acids.  One  of  these  ar^ 

H 

i 

rangements  is  this,  O=S-O-H,  in  which  the  sulphur  is 

II 
O 

O 

n 

represented  as  sexivalent ;  the  other  is  this,  H-O-S-O-H, 
in  which  the  sulphur  is  represented  as  quadrivalent. 
The  facts  that  sulphurous  acid  is  formed  so  readily  by 
simple  contact  of  sulphur  dioxide  with  water ;  that  it 
breaks  down  as  readily  into  sulphur  dioxide  and  water ; 
and  that  it  takes  up  oxygen  and  sulphur  so  readily  to 
form  sulphuric  and  thiosulphuric  acids,  seem  to  speak  in 
favor  of  the  second  of  the  above  formulas.  But,  on  the 


CONSTITUTION  OF  THE  ACIDS  OF  SULPHUR.        229 

other  hand,  certain  facts  established  in  the  study  of  the 
organic  derivatives  of  sulphurous  acid  form  a  strong  ar- 
gument in  favor  of  the  first  formula.  It  is  possible  by 
means  of  reactions  with  certain  compounds  of  carbon  to 
replace  one  of  the  metal  atoms  in  a  sulphite  in  such  a  way 

(E 

as  to  produce  a  compound  of  the  formula  S<  ONa,  the 

conduct  of  which  is  such  as  to  show  that  in  it  the  group 
B  is  in  direct  combination  with  sulphur.  As  the  com- 
pound is  formed  apparently  by  direct  replacement  of  a 
sodium  atom  in  the  sulphite  it  appears  that  this  sodium 
atom  was  in  combination  with  sulphur.  Further,  when 
the  second  sodium  is  replaced  by  the  group  B  a  com- 

(K 

pound   of  the  formula  S-<  OB  is  obtained,  in  which  it 

(o, 

appears  that  one  of  the  groups  is  in  combination  with 
sulphur  and  the  other  with  oxygen.  Now,  it  is  possible 

(01 
by  starting  with  the  compound  S  -<  01,  which  is  made,  as 

(o 

we  shall  see,  by  replacing  one  oxygen  atom  of  sulphui 
dioxide  by  two  chlorine  atoms,  to  introduce  in  the  place 
of  the  two  chlorine  atoms  two  groups,  OB,  and  thus  ob- 

(OB 

tain  the  compound  S-<  OB,  which  plainly  has  the  same 

(o 

composition   as   the   one   represented    by   the   formula 

(E 

S-<  OB,  but  it  has  a  different  constitution.    It  was  found 

(o, 

by  experiment  that  the  two  compounds  have  different 
properties,  and  it  seems  probable,  therefore,  that  the  two 
formulas  given  represent  the  structure  of  the  two  com- 
pounds. As  the  one  which  has  one  group  B  in  combina- 
tion with  sulphur  is  obtained  from  sodium  sulphite  by 
replacement  of  sodium,  the  conclusion  seems  to  be  justi- 
fied that  sulphurous  acid  has  the  constitution  represented 


230  INORGANIC  CHEMISTRY. 

H  f  TT 

i  j  -n- 

by  the  formula  O=S=O  or  SK  OH,  which  may  also  be 

6          i°* 
i 

TT 

written  O2S  <QTT'    At  the  same  time  it  appears  quite  pos- 

sible that  a  sulphurous  acid  of  the  other  constitution, 

OTT 

jj,  maybe  found,  as  there  are  innumerable  ex- 


amples furnished  by  chemistry  of  the  existence  of  two 
or  more  compounds  of  the  same  percentage  composition 
but  different  constitution.  Two  or  more  substances  hav- 
ing the  same  composition  but  different  constitution  are 
said  to  be  isomeric.  Although  examples  of  isomerism 
are  rare  among  the  compounds  of  most  elements,  yet 
among  the  compounds  of  carbon  they  are  met  with  in 
large  numbers,  and  in  studying  these  compounds  a  great 
deal  of  attention  has  been  paid  to  the  phenomena  of 
isomerism.  It  is  possible  that  the  second  form  of  sul- 
phurous acid  cannot  exist  on  account  of  the  tendency  of 
sulphur  to  act  as  a  sexivalent  element.  This  is,  how- 
ever, mere  speculation,  and,  unless  the  suggestion  can  be 
tested  experimentally,  it  is  of  very  little  value.  If  sul- 
phurous acid  has  the  constitution  above  assigned  to  it 
then  the  transformation  of  sulphurous  acid  into  sulphuric 
acid  is  not  simply  a  direct  combination  of  oxygen  with 
sulphur,  but  the  act  must  involve  a  partial  breaking  down 
of  the  sulphurous  acid  and  a  recombination  of  the  con- 
stituents thus  : 

O  O 

H-O-S-H  +  O  =  H-O-S-O-H. 

ll  I'l 

O  O 

If  this  view  is  correct  the  oxygen  displaces  the  hydrogen 
and  then  combines  with  it  and  the  sulphur. 

If  the  action  takes  place  in  the  same  way  with  sulphur 
it  must  be  represented  thus  : 


CONSTITUTION  OF  THE  ACIDS  OF  SULPHUR.        231 

o  o 

H-O-S-H  +  S  =  H-O-S-S-H ; 

II  II 

O  O 

o 

II 
and,  according  to  this,  the  formula  H-O-S-S-H  repre- 

II 
O 

sents  the  structure  of  thiosulphuric  acid.  The  same  con- 
clusion regarding  the  structure  of  thiosulphuric  acid  is 
reached  by  a  consideration  of  one  of  the  methods  by 
which  its  sodium  salt  is  made.  This  is  the  method  which 
consists  in  treating  a  mixture  of  sodium  sulphide  and 
sodium  sulphite  with  iodine.  The  iodine  simply  extracts 
two  atoms  of  sodium  from  the  two  molecules,  and  it 
seems  probable  that  the  residues  of  the  two  molecules 
unite.  If  this  is  correct  the  following  equation  represents 
what  takes  place : 

Q  .Na 

5<Na  S-Na 

Na  +2I  =  0 

O=S-0-Na 

II  O 

O 

and  the  constitution  of  thiosulphuric  acid  is  represented 
S-H 

by  the  formula  O=S-O-H,  which  is  identical  with  that 

II 
O 

given  above.     This  latter  method,  however,  it  should  be 
remarked,  might  also  be  used  as  an  argument  in  favor  of 
H 

the  constitution  O=S-O-S-H  for  thiosulphuric  acid ;  for 

II 
O 

if,  when  the  iodine  acts  upon  the  sulphite,  it  extracts  the 
sodium  atom  which  is  in  combination  with  oxygen,  then 
the  union  of  the  two  residues  would,  in  all  probability, 


INORGANIC  CHEMISTRY. 

take  place  at  this  point,  and  the  constitution  of  the  acid 
would  be  that  represented  by  the  last  formula  given.  It 
will  be  seen  that  our  knowledge  in  regard  to  the  structure 
of  thiosulphuric  acid  is  very  unsatisfactory. 

The  same  remark  may  be  made  in  regard  to  our  knowl- 
edge of  the  structure  of  the  other  acids  of  sulphur.  The 
method  used  in  making  the  salts  of  tetrathionic  acid  is 
suggestive,  and  if  we  knew  the  structure  of  thiosulphuric 
acid,  we  might  draw  a  conclusion  in  regard  to  that  of 
tetrathionic  acid.  The  method  consists  in  treating  the 
sodium  salt  of  thiosulphuric  acid  with  iodine.  It  appears 
probable  that  of  the  two  formulas  above  given  for  thio- 

O 

II 

sulphuric  acid  this  one,  H-S-S-O-H,  is  to  be  preferred, 

II 
O 

for  the  reason  that  it  is  more  probable  that  when  iodine 
acts  upon  sodium  thiosulphate  it  removes  the  sodium 
atom  which  is  in  combination  with  sulphur,  than  that  it 
removes  the  one  which  is  in  combination  with  oxygen. 
If  this  is  true  the  above  formula  for  thiosulphuric  acid 
follows.  Now,  further,  reasoning  in  the  same  way,  it 
appears  probable  that  when  iodine  acts  upon  sodium 
thiosulphate  it  removes  the  sodium  which  is  in  combina- 
tion with  sulphur,  and  in  this  case  the  formation  of  tetra- 
thionic acid  must  be  represented  in  the  following  way : 

O       O  00 


(O-S- 

6 

Na  +  Na 
I         I 

-S-&-O-Na 

6 

=  NaO-^-S-S-S-O-Na  + 

According  to  this,  in  tetrathionic  acid  there  are  four 
sulphur  atoms  in  combination  by  one  affinity  each. 
This  compound  is,  however,  comparatively  stable,  while 
thiosulphuric  acid,  in  which  a  similar  combination  of 
only  two  sulphur  atoms  is  assumed,  is  extremely  un- 
stable. What  value  to  attach  to  such  considerations  as 
these  we  do  not  know. 

Compounds  of  Sulphur  with  Oxygen. — Sulphur,  as  has 
been  repeatedly  stated,  combines  with  oxygen  in  two  pro- 


SULPHUR  DIOXIDE.  233 

portions,  forming  the  oxides  SO2  and  SOS,  or  sulphur  di- 
oxide and  sulphur  trioxide.  Besides  these  two  it  also 
forms  a  sesquioxide,  S2O3,  and  a  heptoxide,  SaO7,  but  com- 
paratively little  is  known  in  regard  to  the  last  two.  The 
one  best  known  is  sulphur  dioxide. 

Sulphur  Dioxide,  SO2. — This,  as  has  been  seen,  is 
formed  when  sulphur  is  burned  in  the  air  or  in  oxygen ; 
and  It  is  also  easily  formed  by  reduction  of  the  higher 
oxides  and  acids  of  sulphur.  Owing  to  the  fact  that 
with  water  it  forms  sulphurous  acid,  it  is  frequently 
called  sulphurous  anhydride.  The  methods  for  making 
it  were  relerred  to  under  sulphuric  acid,  and  the  reac- 
tions involved  were  discussed  then  with  a  sufficient  degree 
of  fulness.  It  need  only  be  said  here  that  in  the  labora- 
tory the  methods  most  commonly  employed  are  :  (1)  Heat- 
ing sulphuric  acid  with  copper ;  (2)  heating  the  acid  with 
carbon  (charcoal)  ;  and  (3)  heating  the  acid  with  sulphur. 
When  carbon  is  used  two  gases  are  formed,  viz.,  carbon 
dioxide  and  sulphur  dioxide  : 

2H2SO4  +  C  =  CO3  +  2SO2  +  2H,O. 

The  gas  can  also  be  made  by  heating  a  mixture  of  a 
metallic  oxide  and  sulphur.  Thus,  when  cupric  oxide  and 
sulphur  are  heated  together  this  reaction  takes  place  : 

2CuO  +  28  =  Cu2S  +  SO2. 

Sulphur  dioxide  is  a  colorless,  transparent  gas,  which 
has  a  pungent,  suffocating  odor,  familiar  as  the  odor  of 
burning  sulphur  matches.  It  is  poisonous,  causing  death 
when  inhaled  in  any  quantity,  and  giving  rise  to  bad 
symptoms  even  in  comparatively  small  quantities.  It 
does  not  readily  give  up  its  oxygen,  so  that  burning 
bodies  are  extinguished  when  introduced  into  it.  It  acts 
something  like  water  in  this  respect.  It  is  more  than 
twice  as  heavy  as  air,  its  specific  gravity  being  2.26. 
When  sulphur  is  burned  in  oxygen  gas  the  sulphur  di- 
oxide formed  occupies  the  same  volume  as  the  oxygen 
used  up,  so  that  there  is  no  change  in  the  volume.  This 


234  INORGANIC  CHEMISTRY. 

can  be  shown  by  the  experiment  here  described.     In  a 
bent  glass  tube,  of  the  form  shown  in  Fig.  8,  there  is 
placed  a  piece  of  sulphur,  and  the 
tube  is  then  half  filled  with  pure 
oxygen    over    mercury.      On    now 
heating  the  tube  at  the  part  where 
the  sulphur  is,  this  burns   and  is 
converted     into     sulphur    dioxide. 
After  the  tube  has  cooled  down  to 
FIG.  s.  the  ordinary  temperature  the  gas  is 

found  to  occupy  the  same  volume  as  before.  This  will 
be  readily  understood  by  the  aid  of  the  following  con- 
siderations :  In  the  reaction 


one  molecule  of  sulphur  dioxide  is  formed  for  every 
molecule  of  oxygen  used  up.  But  a  molecule  of  sulphur 
dioxide  in  the  form  of  gas  occupies  the  same  space  as  a 
molecule  of  oxygen,  so  that,  as  the  space  occupied  by  the 
sulphur  in  the  experiment  is  insignificant,  there  is  no 
change  in  volume  occasioned  by  the  above  reaction. 

Sulphur  dioxide  dissolves  in  water,  as  we  have  seen, 
and  forms  a  liquid  in  which,  judging  by  its  conduct,  sul- 
phurous acid  is  present. 

The  gas  is  easily  liquefied  by  cold  alone.  It  is  only 
necessary  for  this  purpose  to  pass  the  dry  gas  through  a 
tube  surrounded  by  a  freezing  mixture  of  ice  and  salt. 
The  liquid  changes  rapidly  into  gas  under  ordinary  pres- 
sure at  the  ordinary  temperature:  In  this  change  so 
much  heat  is  absorbed  that  a  temperature  of  about  —  60° 
can  be  produced  by  means  of  it  ;  and  a  portion  of  the 
liquid  can  be  solidified. 

Sulphur  dioxide  is  very  stable.  If  heated  to  1200° 
under  pressure,  however,  it  breaks  down  into  sulphur 
trioxide  and  sulphur  : 

'*  3SO2  =  2SO3  +  S. 

When  conducted  into  solutions  of  bases  or  of  carbonates 
the  corresponding  sulphites  are  formed  : 


SULPHUR  TRIOXIDE.  235 

2KOH  +  SO2  =  K2SO3  +  H2O  ; 
K2C03  +  SOi 


Under  certain  conditions,  as  when  the  gas  is  passed 
into  a  hot  solution  of  an  alkali  carbonate,  a  salt  of  the 
general  formula  M2S2O5  is  formed.  This  bears  to  the 
sulphite  the  same  relation  that  the  pyrosulphate  bears 
to  the  sulphate  : 


and  2KHS04  =  K2S2O7  +  H2O. 


Sulphur  dioxide  is  used  extensively  for  the  purpose  of 
bleaching  silk,  wool,  straw,  and  basket-ware.  In  order 
that  it  may  bleach,  however,  water  must  be  present,  so 
that  it  appears  that  the  true  bleaching  agent  in  this  case 
is  sulphurous  acid  and  not  the  dioxide.  When  we  con- 
sider that  sulphur  dioxide  does  not  readily  take  up  nor 
give  up  oxygen,  while  sulphurous  acid  does  readily  take 
it  up,  the  necessity  of  having  water  present  in  the 
bleaching  process  at  once  becomes  apparent.  The 
bleaching  in  some  cases  certainly  consists  in  abstracting 
oxygen  from  the  colored  substances,  and  thus  converting 
them  into  colorless  products.  In  other  cases  it  is  due  to 
the  formation  of  compounds  of  sulphurous  acid  with  the 
dye-stuffs. 

Sulphur  dioxide  is  not  only  a  bleaching  agent  like 
chlorine,  but  like  chlorine  it  is  also  a  disinfectant.  It 
has  to  some  extent  the  power  to  destroy  the  organisms 
which  cause  changes  in  organic  substances.  It  prevents 
fermentation  and  is  therefore  used  as  a  preservative.  Its 
power  to  destroy  the  germs  of  disease,  that  is,  to  disinfect, 
is  not  as  great  as  is  frequently  supposed.  Much  larger 
quantities  are  necessary  for  this  purpose  than  are  com- 
monly used. 

Sulphur  Trioxide,  SO3.  —  This  compound  is  made  by 
passing  sulphur  dioxide  and  oxygen  together  over  heated 
platinum  in  a  finely  divided  state.  It  is  obtained  most 
readily  by  heating  disulphuric  acid,  which  breaks  up 


236  INORGANIC  CHEMISTRY. 

easily  into  sulphur  trioxide  and  ordinary  sulphuric  acid 
according  to  the  equation 

HaSa07  =  HaS04  +  S03. 

Similarly,  the  acid  sulphates  of  the  alkali  metals  yield 
the  corresponding  normal  sulphates  and  sulphur  trioxide  : 

2NaHS04  =  Na2SO4  +  SO3  +  H2O. 

It  is  now  manufactured  on  the  large  scale  by  passing  sul- 
phur dioxide  and  oxygen  together  over  asbestos  covered 
with  finely  divided  platinum,  and  the  product  thus  ob- 
tained is  passed  into  ordinary  sulphuric  acid  for  the  pur- 
pose of  making  "  solid  sulphuric  acid "  which,  as  has 
been  stated,  is  almost  pure  disulphuric  acid,  H2S2O7. 

Sulphur  trioxide  is  a  white  crystallized  solid  which 
appears  to  exist  in  two  modifications.  The  one  is  a  liquid 
at  ordinary  temperatures,  but  it  solidifies  at  about  16°. 
According  to  the  latest  investigations  there  is  but  one 
modification  of  the  oxide.  It  is  a  solid  which  melts  at 
14.8°,  forming  a  liquid  which  boils  at  46°.  In  contact 
with  the  air  the  oxide  gives  off  thick  fumes  which  are 
partly  due  to  the  great  power  of  the  compound  to  com- 
bine with  water.  Water  acts  with  violence  upon  it,  the 
heat  evolved  in  the  act  being  39,170  cal.  It  also  acts 
upon  substances  containing  hydrogen  and  oxygen  in  much 
the  same  way  that  concentrated  sulphuric  acid  does, 
charring  them  by  abstracting  the  hydrogen  and  oxygen. 
It  acts,  however,  more  violently  in  this  way  than  sulphuric 
acid  does.  With  water  it  forms  sulphuric  acid,  and  it  is, 
therefore,  called  sulphuric  anhydride.  The  reaction  in- 
volved in  passing  from  sulphur  trioxide  to  sulphuric  acid 
is  of  a  kind  which,  as  we  have  seen,  is  frequently  met  with 
both  with  acidic  oxides  or  anhydrides,  and  with  basic  or 
metallic  oxides ;  and  it  is  desirable  that  it  should  here 
be  studied  a  little  more  carefully  than  it  has  yet  been. 
What  we  know  is  that  when  sulphur  trioxide  acts  upon 
water  there  is  a  great  deal  of  heat  evolved,  and  com- 
pounds of  different  composition  are  obtained.  The  com- 
position of  these  compounds  is  represented  by  the  formu- 


ACIDIC  OXIDES  AND    WATER.  237 

las  SO3  +  H2O,  SO3  +  2H2O,  and  SO3  +  3H2O.  So,  too, 
when  calcium  oxide  or  lime,  CaO,  acts  upon  water,  there 
is  great  evolution  of  heat,  and  a  compound  is  formed 
the  composition  of  which  is  represented  by  the  formula 
CaO  4-  H2O.  But  these  formulas  do  not  attempt  to  give 
any  account  of  what  takes  place  in  the  chemical  acts  re- 
ferred to.  That  the  water  is  not  present  in  the  com- 
pounds as  water  seems  evident,  in  the  first  place  from 
the  conduct  of  the  compounds,  and  in  the  second  place 
from  the  amount  of  heat  evolved  in  the  act  of  combina- 
tion. Now,  taking  the  chemical  conduct  of  the  substances 
into  consideration,  they  appear  to  contain  hydrogen  in 
•combination  with  oxygen,  and  their  conduct  becomes 
comprehensible  on  the  supposition  that  the  group  known 
as  hydroxyl,  (-O-H),  is  present.  This  view  has  been 
found  to  be  in  accordance  with  a  large  number  of  facts, 
and  it  is  of  great  assistance  in  dealing  with  these  facts. 
The  view  is  distinctly  this  :  When  an  acidic  oxide  acts 
upon  water  it  is  converted  into  a  hydroxyl  compound 
which  has  acid  properties,  as  shown  in  this  equation  : 

roH 
0        TT  o         OH 

II  H,U 

O^S^O  +  H20  =  S 

TT  r\ 

H20 

LOH 

According  to  this,  each  molecule  of  water  is  decomposed 
and  each  hydrogen  atom  in  the  resulting  compound 
is  in  combination  with  oxygen.  The  same  kind  of  action 
takes  place  in  the  case  of  some  basic  oxides,  as  shown, 
for  example,  in  the  case  of  calcium  oxide  : 


By  means  of  certain  reagents  which  will  be  taken  up 
later  it  is  possible  to  replace  oxygen  in  compounds  by 
chlorine,  two  chlorine  atoms  taking  the  place  of  one 
oxygen  atom.  So,  also,  when  the  oxygen  of  the  hydrox- 
ides is  replaced  by  chlorine,  the  result  is  that  each 


238  INORGANIC  CHEMISTRY. 

hydroxyl  group  is  replaced  by  one  atom  of  chlorine.  This 
is  easily  understood.  For,  if  in  calcium  hydroxide  each 
oxygen  atom  should  be  replaced  by  two  chlorine  atoms, 
the  result  would  be  a  combination  of  atoms  represented 

C1C1  -H 
thus,  Ca<cici-H  »   an^   fr°m   ^s>  hydrochloric   acid 

would  be  given  off,  leaving  calcium  chloride  : 

Cl 


O—  TT 

The  result,  therefore,  of  treating  the  hydroxide  C/a<Q_Tj- 

in  such  a  way  as  to  replace  each  of  the  two  oxygen  atoms 
by  two  chlorine  atoms  is  the  formation  of  a  compound 

Cl 
Ca<™,  which  is  derived  from  the  hydroxide  by  replacing 

each  hydroxyl  group  by  one  atom  of  chlorine.  These 
facts  lend  support  to  the  hydroxyl  hypothesis. 

Just  as  sulphur  trioxide  acts  upon  water  to  form  sul- 
phuric acid,  so  it  acts  upon  metallic  oxides,  forming 
sulphates.  Thus  with  calcium  oxide  it  forms  calcium 
sulphate  : 

CaO  +  SO3  =  CaSO4. 

It  combines  also  with  hydrochloric  acid,  forming  the 
compound  SO3HC1.  The  action  in  this  case  is  analogous 
to  that  which  takes  place  with  water,  and  is  represented 
thus  : 

O  O 

O=S=0  +  H-C1  =  0-S-O-H. 

I 
Cl 

The  analogy  between  this  and  the  action  with  water  will 
be  apparent  by  a  consideration  of  the  following  equation  : 

O  O 

O=S=O  +  H-O-H  =  O=S-O-H. 

I 

O 

I 

H 


ACID  CHLORIDES  OF  SULPHUR.  239 

The  chlorine  compound  appears  to  be  ordinary  sul- 
phuric acid  in  which  one  hydroxyl  group  has  been  replaced 
by  a  chlorine  atom.  When  the  compound  is  treated  with 
water  hydrochloric  acid  is  given  off  and  sulphuric  acid 
is  formed,  thus  : 

(  01  (  OH 

S^  OH  +  HOH  =  S-J  OH  +  HC1. 
(O,  (O, 

Acid  Chlorides  of  Sulphur.  —  In  the  introduction  to 
this  chapter  reference  was  made  to  certain  compounds 
which  sulphur  forms  with  oxygen  and  chlorine,  known 
as  acid  chlorides,  because  when  brought  in  contact  with 
water  they  form  acids.  Strictly  speaking,  all  the  com- 
pounds of  sulphur  and  chlorine  are  of  this  order  ;  the 
name  is,  however,  generally  applied  to  those  compounds 
which  are  derived  from  the  acids  by  replacement  of  the 
hydroxyl  by  chlorine.  Thus  the  acid  chloride  of  sul- 

01 
phuric  acid  is  SO2<Cj,  which  is  plainly  sulphuric  acid  in 

which  the  two  hydroxyl  groups  have  been  replaced  by 
chlorine.     So,  also,  the  compound  SO<™  is  related  in 

the   same   way   to   a   sulphurous   acid   of    the   formula 

OH 

,  which,  however,  does  not  appear  to  be  ordi- 


nary sulphurous  acid.     But  the  existence  of  the  com- 

Pl  PI 

pounds  SO2<^  and  SO<QJ,  which  are  to  be  regarded 

respectively  as  sulphuric  and  sulphurous  acids  in  which 
both  hydroxyls  are  replaced  by  chlorine,  suggests  the 

Cl  PI 

existence  of  the  two  compounds  SO2<         and 


derived  from  the  acids  by  the  replacement  of  one  hy- 
droxyl in  each  by  chlorine.  The  former  compound  is 
known,  the  latter  is  not. 

Thionyl  Chloride,  SOC12,  can  be  made  from  sulphur 
dioxide  by  replacing  one  oxygen  atom  by  two  chlorine 
atoms.  This  is  effected  by  treating  sulphur  dioxide  with 
phosphorus  pentachloride,  PC15,  a  compound  which 
readily  gives  up  a  part  or  all  of  its  chlorine  and  takes  up 


240  INORGANIC  CHEMISTRY. 

oxygen  in  its  place,  and  is  therefore  extensively  used  in 
chemistry  for  the  purpose  of  replacing  oxygen  by  chlo- 
rine. With  sulphur  dioxide  the  reaction  takes  place 
thus: 

SO3  +  PC16  =  SOC12  +  POC13. 

The  compound  is  also  formed  by  treating  sodium  sul- 
phite with  phosphorus  pentachloride,  when  this  reaction 
takes  place  : 

SO(ONa)2  +  2PC15  =  SOC12  +  2NaCl  +  2POC13. 

In  order  to  understand  this  reaction  it  appears  neces- 
sary that  sodium  sulphite  should  be  represented  by  the 
formula  given,  which,  as  will  be  seen,  is  not  in  accordance 
with  the  conclusion  reached  regarding  the  structure  of 
sulphurous  acid.  It  may  be  that  the  reaction  is  more 
complicated  than  here  represented,  and  that  sulphur 
dioxide  is  first  given  off,  and  that  this  then  acts  upon  the 
pentachloride. 

Thionyl  chloride  is  a  liquid  with  a  strong  characteristic 
odor.  It  acts  readily  upon  water,  forming  sulphur  dioxide 
and  hydrochloric  acid  according  to  the  equation 

SOC12  +  H2O  =  SO2  +  2HC1. 

Sulphuryl  Chloride,  SO2C12»  is  formed  by  the  direct 
action  of  chlorine  upon  sulphur  dioxide  in  the  sunlight, 
the  action  being  similar  to  that  which  takes  place  when 
sulphur  dioxide  and  oxygen  unite  to  form  the  trioxide  : 


SO2  +  0    =  SO3. 

01 
It  is  best  prepared  by  heating  the  compound  SOa 

to  170°-180°,  when  the  following  reaction  takes  place  : 

2SO«-<OH  =  H*SO<  +  SO*C1*' 

It  is  a  liquid  which  is  easily  decomposed  by  water,  as 
represented  in  the  equation 

S03C12  +  2H20  =  S02(OH)2  +  2HC1. 


OXYGEN  COMPO  UNDS  OF  SELENIUM  AND  TELL  URIUM.  241 
With  half  the  quantity  of  water  required  to  effect 
the  above  decomposition  cJdorsulphuric  acid,  S0a</yrr, 
is  formed  : 

so*<ci +H»°  =  8°«<oa  +  HCL 

Chlorsulphuric  Acid,  or  Sulphuryl-hydroxyl  Chloride, 
SOa<QH,  is  also  formed,  as  has  been  shown,  by  the  direct 
action  of  hydrochloric  acid  upon  sulphur  trioxide : 


and,  further,  by  the  action  of  phosphorus  pentachloride 
upon  sulphuric  acid  : 

SO°<OH  +  PC1»  =  8°.<OH  +  POC1°  +  HCL 

Like  sulphuryl  chloride  it  is  decomposed  by  water,  yield- 
ing hydrochloric  acid  and  sulphuric  acid. 

Compounds  of  Selenium  and  Tellurium  with  Oxygen  and 
with  Oxygen  and  Hydrogen. — For  the  purposes  of  this 
book  it  is  not  necessary  to  go  into  details  in  regard  to 
the  compounds  of  selenium  and  tellurium.  In  the  intro- 
duction to  this  chapter  it  was  stated  that  selenium  and 
tellurium  form  compounds  with  oxygen  corresponding 
to  sulphur  dioxide  and  sulphur  trioxide.  Both  of  these 
oxides  of  tellurium  are  known,  together  with  a  third  of 
the  composition  represented  by  the  formula,  TeO,  while 
only  one  oxide  of  selenium,  the  dioxide,  SeO2,  is  known. 
On  the  other  hand,  the  acids  of  selenium  and  tellurium, 
corresponding  to  sulphurous  and  sulphuric  acids,  are 
known. 

Selenious  Acid,  H2SeO3. — This  compound  is  formed  by 
oxidation  of  selenium  in  presence  of  water  or  by  dissolv- 
ing selenium  dioxide  in  water.  It  is  a  crystallized  solid, 
which  attracts  moisture  from  the  air,  and  when  heated 
breaks  down  into  selenium  dioxide  and  water.  It  forms 
two  series  of  salts,  the  acid  selenites,  MHSeO3,  and  the  nor- 


242  INORGANIC  CHEMISTRY. 

mal  selenites,  MaSeO3.  While  in  composition  the  acid 
and  its  salts  are  analogous  to  sulphurous  acid  and  the 
sulphites,  in  conduct  there  is  a  marked  difference.  Sul- 
phurous acid  tends,  as  we  have  seen,  to  take  up  more 
oxygen  and  form  sulphuric  acid,  while  selenious  acid 
gives  up  its  oxygen  with  great  ease  and  yields  selenium. 
It  is  evident  that  the  affinity  of  selenium  for  oxygen  is 
much  less  than  that  of  sulphur  for  oxygen.  When  a 
solution  of  selenious  acid  is  treated  with  sulphur  dioxide 
the  acid  is  reduced  to  selenium  : 

HaSe03  +  2SO2  +  H20  =  2H2SO4  +  Se. 

The  same  method  of  investigation  which  leads  to  the 
conclusion  that  sulphurous  acid  has  the  formula 

TT 

SOa<Q-rr,  when  applied  to  selenious  acid,  leads  to  the 

OH 

conclusion  that  this  has  the  constitution  SeO<Q-rr. 

Selenic  Acid,  H2SeO4,  is  formed  by  the  action  of  power- 
ful oxidizing  agents  like  saltpeter,  or  chlorine  or  bromine 
in  water  solution  on  selenium.  The  action  with  chlorine 
is  represented  thus : 

Se  +  601  +  4H2O  =  H2SeO4  +  6HC1. 

It  is  a  liquid  resembling  sulphuric  acid.  It  cannot,  how- 
ever, be  obtained  in  pure  condition  on  account  of  its  ten- 
dency to  give  up  oxygen  and  pass  over  into  selenious 
acid,  in  which  respect  it  plainly  differs  markedly  from 
sulphuric  acid.  It  gives  up  its  oxygen  to  other  substances 
much  more  readily  than  sulphuric  acid  does.  This  is 
seen  in  its  action  upon  hydrochloric  acid,  which  takes 
place  as  represented  in  the  equation 

H2Se04  +  2HC1  =  H2SeO3  +  01,  +  H2O. 

Sulphuric  acid  does  not  act  upon  hydrochloric  acid, 
but  it  does  act  upon  hydrobromic  and  hydriodic  acids, 
and  its  action  upon  hydrobromic  acid  is  very  similar  to 


ACID  CHLORIDES  OF  SELENIUM.  243 

that  which  takes  place  between  selenic  acid  and  hydro- 
chloric acid  : 

H3S04  +  2HBr  =  SO2  +  Br2  +  2H2O  ; 

the  only  essential  difference  is  that  the  sulphurous  acid 
breaks  down  into  sulphur  dioxide  and  water,  while  the 
selenious  acid  remains  in  solution  as  such. 

Selenium  Dioxide,  SeO2.  —  This  analogue  of  sulphur 
dioxide  is  made  by  burning  selenium,  or  by  heating 
selenious  acid.  It  is  a  solid  crystallized  substance 
which  can  be  sublimed  without  undergoing  decompo- 
sition and  without  melting.  As  we  have  seen,  it  dissolves 
readily  in  water,  forming  selenious  acid.  It  loses  its  oxy- 
gen readily  in  contact  with  substances  which  have  the 
power  to  unite  with  oxygen,  and  it  is  thus  reduced  to 
selenium.  Hence  when  it  is  sublimed  particles  of  dust 
which  may  be  present  in  the  vessel  effect  partial  reduction, 
and  the  product,  instead  of  being  white,  as  the  oxide  is 
when  pure,  is  colored  by  the  particles  of  selenium. 

Acid  Chlorides  of  Selenium.  —  Selenium  combines  with 
oxygen  and  chlorine,  forming  compounds  analogous  to  the 
acid  chlorides  of  sulphur.  The  principal  one  of  these  is 
sdenyl  chloride,  SeOCl2,  which  is  formed  by  bringing  to- 
gether selenium  tetrachloride  and  selenium  dioxide  : 

SeCl4  +  SeO2  =  2SeOCl2. 

By  water  it  is  decomposed,  forming  selenious  and  hydro- 
chloric acids  : 


Just  as  sulphur  trioxide  combines  with  hydrochloric 
acid,  so  also  does  selenium  dioxide.  It  forms  products 
which  are  represented  by  the  formulas  SeO2.2HCl  and 
SeO2.4HCl.  It  is  quite  possible  that  the  former  is  a 

Cl>Se<OH 
chloride  of  selenious  acid  of  the  constitution  Cl  OH  ' 


244  INORGANIC  CHEMISTRY. 

and  that  the  second  should  be  represented  by  the  for. 
mula -rr_/ni2\>Se<QTT,  a  few  compounds  of  this  kind 

being  known,  as  will  be  pointed  out. 

Tellurious  Acid,  H2TeO3. — Tellurious  acid  is  formed  by 
treating  tellurium  tetrachloride  with  water.  It  is  possible 
that  the  first  action  causes  the  formation  of  normal  tel- 
lurious  acid,  Te(OH)4,  and  that  this  then  breaks  down 
into  tellurious  acid,  H2TeO3,  and  water : 


Cl  HOH             f  OH 

T    ,  Cl  ,    HOH_I 

le^  Cl  +HOH  -  1 

Cl  HOH             L  OH 

(OH 
Te^g|  =  Te]oH  +  H20. 

OH  ^° 


The  potassium  salt  of  the  acid  is  formed  by  melting 
together  tellurium  dioxide  and  potassium  carbonate : 

K2CO3  +  TeO2  =  K2TeO3  +  CO,. 

If  the  salt  thus  formed  is  dissolved  in  water  and  nitric 
acid  added  to  the  solution,  tellurious  acid  is  thrown  down: 

K2TeO3  +  2HNO3  =  2KNO3  +  H2TeO3. 

It  is  a  solid  which  easily  loses  water  and  is  thus  trans- 
formed into  tellurium  dioxide. 

Telluric  Acid,  H2TeO4. — This  acid  is  formed  by  melting 
tellurious  acid  with  saltpeter  and  other  oxidizing  agents. 
When  the  solution  of  the  acid  is  evaporated  to  crystal- 
lization the  solid  compound  deposited  has  the  compo- 
sition H6TeO6,  and,  according  to  what  was  learned  in 
studying  sulphuric  acid,  it  appears  probable  that  this  is 
normal  telluric  acid,  Te(OH)6.  When  normal  telluric 
acid  is  heated  to  a  little  above  100°  it  loses  water  and  is 


OXIDES  OF  TELLURIUM.  245 

transformed  into  the  acid  H2TeO4,  corresponding  to  or- 
dinary sulphuric  acid,  from  which  most  of  the  tellurates 
are  derived : 

Te(OH)6  =  Te02(OH)2  +  2H2O. 

Heated  higher,  to  about  160°,  the  acid  is  decomposed 
into  tellurium  trioxide  and  water : 

Te02(OH)a  =  TeO3  +  H2O. 

Although  most  of  the  tellurates  are  simple  salts  of  the 
acid  H2TeO4,  others  are  derived  from  more  complex  forms 
of  the  acid.  One  of  these  is  analogous  to  disulphuric  acid. 
Another  is  derived  from  four  molecules  of  telluric  acid, 
Te(OH)6,  by  loss  of  eleven  molecules  of  water. 

Oxides  of  Tellurium. —  Tellurium  monoxide  is  formed  by 
heating  sulphur  trioxide  and  tellurium  together  in  a 
vacuum.  Tellurium  dioxide  is  formed  by  burning  tellu- 
rium or  by  oxidizing  it  with  nitric  acid ;  and,  further,  by 
the  decomposition  of  tellurious  acid  by  heat.  It  crystal- 
lizes and  is  but  slightly  soluble  in  water. 

The  trioxide,  TeO3,  is  formed  by  heating  telluric  acid 
to  a  high  temperature.  Its  conduct  is  entirely  different 
from  that  of  sulphur  trioxide.  While  the  latter  acts  with 
violence  upon  water,  and  readily  upon  metallic  oxides 
and  hydrochloric  acid,  the  former  does  not  act  readily 
upon  any  of  these  substances.  It  is  insoluble  in  hot  as 
well  as  cold  water. 

Sulphotelluric  Acid  is  an  example  of  the  sulphur  acids 
referred  to  on  p.  141.  While  the  acid  itself  is  not  known, 
a  potassium  salt  of  the  formula  K2TeS4  is  known.  This 
is  plainly  analogous  to  the  salt  of  the  oxygen  acid, 
K2TeO4,  differing  from  it  only  in  containing  sulphur  in 
place  of  the  oxygen. 

FAMILY  VI,  GBOUP  A. 

Group  A,  Family  VI,  includes  chromium,  molyb- 
denum, tungsten,  and  uranium.  All  of  these  show 
some  resemblance  to  the  elements  of  the  sulphur  group, 


246  INORGANIC  CHEMISTRY. 

but  they  also  appear  in  entirely  different  characters, 
forming  compounds  of  a  kind  unknown  among  the 
derivatives  of  sulphur  and  its  analogues.  The  relation 
which  these  elements  bear  to  sulphur  is  much  like  the 
relation  which  manganese  bears  to  chlorine.  The  re- 
semblance to  sulphur  is  seen  mainly  in  the  formation  of 
acids  of  the  formulas  H2CrO4,  H2MoO4,  H2WO4,  and 
H2UO4 ;  and  the  oxides  CrO3,  MoO3,  WO3,  and  UO3.  Most 
of  these  acids  yield  complicated  derivatives,  all  of  which 
can,  however,  be  explained  by  the  same  method  as  that 
used  in  the  case  of  periodic  acid.  The  common  salts  of 
chromic  acid  are  derived  from  dichromic  acid,  which  is 
analogous  to  disulphuric  acid.  They  have  the  general 
formula  M2Cr2O7.  So,  too,  salts  of  molybdic  acid  are 
known  which  are  derived  from  the  simple  form  of  the 
acid,  H2MoO4,  and  others  which  are  derived  from  a  dimo- 
lybdic  acid,  H2Mo2O7,  and  from  more  complicated  forms. 
Tungsten  has  a  wonderful  power  of  forming  complex 
acids.  All  of  them,  however,  can  be  referred  to  the 
simple  form  H2WO4.  And,  finally,  uranic  acid  forms  salts 
which  for  the  most  part  are  derived  from  diuranic  acid, 
H2U2O7.  All  the  most  important  of  these  compounds 
will  be  taken  up  later. 

When  the  acids  of  chromium,  molybdenum,  tungsten, 
and  uranium  lose  oxygen  they  form  compounds  which 
have  little  or  no  acid  character.  The  lower  oxides  of 
chromium  form  salts  with  acids,  and  these  bear  a  general 
resemblance  to  the  salts  of  aluminium,  iron,  and  manga- 
nese. The  chromates  lose  their  oxygen  quite  readily 
when  acids  are  present  with  which  the  chromium  can 
enter  into  combination  in  its  capacity  as  a  base-forming 
element.  Thus,  when  potassium  chromate  K2CrO4  is 
treated  with  hydrochloric  acid  in  the  presence  of  some- 
thing which  can  take  up  oxygen,  decomposition  takes 
place  thus : 

2K2Cr04  +  10HC1  =  4KC1  +  2CrCl3  +  5H2O  +  30. 

With  sulphuric  acid  the  action  takes  place  as  repre- 
sented in  this  equation : 

2K2Cr04  +  5H2S04  =  2K2SO4  +  Cra(S04)3  +  5H20  +  3O. 


COMPOUNDS  OF  CHROMIUM,  MOLYBDENUM,  ETC     247 

In  both  these  cases  the  chromium  enters  into  combina- 
tion as  a  trivalent  base-forming  element,  taking  the  place 
of  three  atoms  of  hydrogen  in  hydrochloric  acid  in  the 
first  case,  and  of  three  atoms  of  hydrogen  in  the  sulphuric 
acid  in  the  second.  Molybdenum  and  tungsten  do  not 
form  salts  of  this  character ;  indeed  they  seem  to  be 
practically  devoid  of  basic  properties.  Uranium,  on  the 
other  hand,  forms  some  curious  salts  which  differ  from 
the  simple  metallic  salts  which  we  commonly  have  to 
deal  with.  These  are  the  so-called  uranyl  salts,  which 
are  regarded  as  acids  in  which  the  hydrogen  is  either 
wholly  or  partly  replaced  by  the  group  UO2,  which  is 
bivalent.  Thus,  the  nitrate  has  the  formula  (UO2)(NO3)2, 
the  sulphate  is  (UO2)SO4,  etc.  These  salts  are  derived 
from  the  compound  UO2(OH)2  acting  as  a  base,  whereas 
this  compound  has  also  distinctly  acid  properties. 


CHAPTER  XV. 

NITROGEN  — THE    AIR  — ARGON. 

NITROGEN,  N  (At.  Wt.  13.93). 

General. — Nitrogen  bears  to  a  group  of  elements  rela- 
tions very  similar  to  those  which  oxygen  bears  to  the  sul- 
phur group,  and  fluorine  to  the  chlorine  group.  There 
are  easily  recognized  resemblances  between  it  and  the 
members  of  the  group,  and  yet  there  are  some  marked 
differences.  As  has  been  stated,  and  as  is  seen  from  its 
position  in  the  periodic  system,  nitrogen  is  trivalent 
towards  hydrogen,  as  shown  in  the  compound  NH3,  while 
it  is  both  trivalent  and  quinquivalent  towards  oxygen,  as 
appears  to  be  shown  in  N2O3  and  N206.  The  principal 
hydrogen  compound,  ammonia,  is  entirely  different  in 
character  from  those  of  chlorine  and  sulphur,  for,  while 
these  are  acid,  ammonia  has  in  a  marked  way  the  char- 
acter of  a  base,  acting,  however,  in  a  peculiar  way  upon 
acids  to  form  salts.  The  two  oxides  above  referred  to 
are  acidic,  forming  the  acids  HNO2  and  HNO3,  which  are 
known  as  nitrous  and  nitric  acids  respectively. 

Occurrence  of  Nitrogen. — It  was  discovered  by  Lavoi- 
sier and  Scheele  towards  the  end  of  the  last  century  that 
the  air  consists  of  two  gases,  one  of  which  is  oxygen,  and 
they  showed  that  when  the  oxygen  is  removed  the  gas 
which  is  left  has  not  the  power  to  support  combustion 
nor  to  support  respiration.  This  gas  was  first  called 
azote  (from  a,  privitive,  and  goon/co?,  life),  and  this  name 
is  still  retained  in  France,  the  symbol  in  use  in  that 
country  being  Az,  whereas  in  all  others  the  symbol  is  N. 
This  is  the  only  case  in  which  there  is  a  difference  of 
usage  in  respect  to  the  symbols  of  the  chemical  elements 
in  different  countries.  The  name  nitrogene  was  given  to 
it  later,  from  the  fact  that  it  is  a  constituent  of  niter  or 
saltpeter,  KNO3  (nitrum,  saltpeter,  and  yeveiv,  to  pro- 

(248) 


PREPARATION  OF  NITROGEN.  249 

duce),  and  this  is  the  origin  of  the  English  name  nitro- 
gen. Not  only  is  nitrogen  found  free  in  the  air,  but  it  is 
found  in  combination  in  a  large  number  of  substances  in 
nature.  It  is  found  in  the  nitrates,  or  salts  of  nitric  acid, 
particularly  as  the  potassium  salt  KNO3,  and  the  sodium 
salt  NaNO3,  which  occurs  in  enormous  quantities  in 
Chili,  and  is  therefore  known  as  Chili  saltpeter.  It  is 
also  found  in  the  form  of  ammonia,  which  is  a  compound 
of  nitrogen  and  hydrogen  of  the  formula  NH3.  Ammonia 
occurs  in  small  quantity  in  the  air,  and  is  formed  under 
a  variety  of  conditions,  to  which  reference  will  be  made 
when  the  substance  is  treated.  Nitrogen  occurs,  further, 
in  combination  in  many  animal  substances. 

Preparation. — The  most  convenient  way  to  prepare 
nitrogen  is  by  burning  in  a  closed  vessel  something  which 
does  not  give  a  gaseous  product  of  combustion ;  or  by 
passing  air  over  something  which  has  the  power  to  unite 
with  oxygen.  The  best  substance  to  use  for  the  first 
purpose  is  phosphorus,  which  burns  readily  and  yields  a 
solid  product,  soluble  in  water.  It  is  only  necessary, 
therefore,  to  place  a  piece  of  phosphorus  in  a  floating 
vessel  on  the  surface  of  water,  set  fire  to  it,  and  immedi- 
ately place  over  it  a  closed  bell-jar.  As  soon  as  the  oxygen 
is  used  up  the  combustion  stops,  and  the  vessel  then 
contains  the  residual  nitrogen,  and  the  walls  are  covered 
with  a  thin  layer  of  phosphorus  pentoxide,  P2O5.  This 
is  soon  converted  by  the  water  into  phosphoric  acid, 
which  dissolves.  Another  convenient  method  for  pre- 
paring nitrogen  consists  in  passing  air  over  copper 
heated  in  a  tube.  The  copper  takes  up  the  oxygen 
readily,  and  the  nitrogen  passes  on.  Another  good 
method  consists  in  exposing  to  the  air  copper  turnings 
partly  covered  with  a  solution  of  ammonia  in  a  vessel  so 
arranged  as  to  allow  free  access  of  air  while  the  escape 
of  the  gas  in  the  vessel  is  prevented.  This  mixture  ab- 
sorbs oxygen  slowly  at  the  ordinary  temperatures. 
Nitrogen  can  also  be  made  from  other  substances  than 
the  air.  Thus,  when  chlorine  is  passed  into  a  water 
solution  of  ammonia  this  reaction  takes  place  : 
NHS  +  3C1  =  N  +  3HC1 ; 


250  INORGANIC  CHEMISTRY. 

but  the  hydrochloric  acid  combines  at  once  with  ammonia 
to  form  ammonium  chloride,  NH4C1  : 

NH3  +  HC1  =  NH4C1  ; 

so  that  the  only  gaseous  product  is  nitrogen.  This  ex- 
periment is  more  or  less  dangerous,  for  if  all  the  ammonia 
should  be  used  up,  and  the  passage  of  chlorine  continued, 
a  compound  of  nitrogen  and  chlorine  which  is  extremely 
explosive  is  formed.  Finally,  nitrogen  can  be  made  by 
heating  ammonium  nitrite,  NH4NO2,  either  dry  or  in  solu- 
tion. The  hydrogen  and  oxygen  of  the  compound  unite 
to  form  water  and  the  nitrogen  is  set  free  : 


The  nitrogen  prepared  from  the  air  is  never  pure,  as 
there  are  always  present  in  the  air  other  substances  be- 
sides nitrogen  and  oxygen  ;  and  while  some  of  these  can 
be  removed  without  serious  difficulty,  others  cannot  be. 
Properties.  —  Nitrogen  is  a  colorless,  tasteless,  inodor- 
ous gas.  It  has  been  converted  into  a  liquid  by  subject- 
ing it  to  a  very  low  temperature  and  high  pressure.  The 
liquid  solidifies  at  —203°.  A  liter  of  nitrogen  under 
standard  conditions  weighs  1.250  grams.  Its  specific 
gravity  (air  —  1)  is  0.971.  It  does  not  support  combus- 
tion, nor  does  it  burn.  This  latter  fact  is  obvious,  for,  if 
nitrogen  had  the  power  to  combine  with  oxygen  when 
the  temperature  of  the  mixture  is  elevated,  it  is  plain 
that  this  process  of  combustion  would  long  ago  have  taken 
place,  leaving  one  or  the  other  of  the  two  gases  and  the 
product  of  combustion  as  the  constituents  of  the  air.  Ni- 
trogen not  only  does  not  combine  with  oxygen  readily,  but 
it  does  not  combine  readily  with  any  other  element  ex- 
cept at  very  high  temperature,  and  then  with  only  a  few. 
Just  as  it  does  not  support  combustion,  so  also  it  does 
not  support  respiration.  Animals  would  die  in  it,  not  on 
account  of  any  active  poisonous  properties  possessed  by 
it,  but  for  lack  of  oxygen.  In  the  air  it  serves  the  useful 
purpose  of  diluting  the  oxygen.  If  the  air  consisted  only 
of  oxygen,  all  processes  of  combustion  would  certainly  be 
much  more  active  than  they  now  are.  What  effect  the 


THE  AIR.  251 

continued  breathing  of  oxygen  would  have  upon  animals 
it  is  impossible  to  say. 

The  Air. — The  atmosphere  of  the  earth,  commonly 
called  the  air,  consists  essentially  of  the  two  elements 
nitrogen  and  oxygen  in  the  proportion  of  79  volumes  of 
nitrogen  to  21  volumes  of  oxygen,  or,  by  weight,  of  77 
per  cent  of  nitrogen  and  23  per  cent  of  oxygen.  Wher- 
ever air  has  been  collected  and  analyzed  it  has  been 
found  to  have  practically  the  same  composition.  Never- 
theless very  accurate  analyses  have  shown  that  the  com- 
position of  the  air  is  subject  to  slight  variations.  To  de- 
cide whether  the  air  is  a  chemical  compound  or  a  me- 
chanical mixture  requires  a  careful  examination  of  a 
number  of  facts.  The  evidence  may  be  summed  up  as 
follows : 

(1)  If  nitrogen  and  oxygen  are  mixed  together  the  mix- 
ture conducts  itself  in  exactly  the  same  way  as  air.    The 
mixing  is   not   attended   by  any  phenomena   indicating 
chemical  action.     Generally  the  chemical  combination  of 
two  elements  is  accompanied  by  an  evolution  of  heat,  and 
whenever   a   chemical   act   takes    place   there   is    some 
change  in  the  temperature    of   the    substances.     When 
nitrogen  and  oxygen  are  brought  together  there  is  no 
change  in  the  temperature  of  the  gases. 

(2)  The  composition  of  a  chemical  compound  is  con- 
stant.    The  law  of  definite  proportions  is  founded  upon 
a  very  large  number  of  observations,  and  in  all  cases  in 
which  we  have  independent  evidence  that  chemical  action 
takes  place  it  is  found  that  the  substances  combine  in 
exactly  the  same  proportions  to  form  the  same  product. 
Variation  in  the  composition  of  a  chemical  compound  is 
not  known.     The  composition  of  the  air  varies  slightly, 
according  to  circumstances,  and  this  fact  may  be  regarded 
as  evidence  that  the  air  is  not  a  chemical  compound. 

(3)  Air  dissolves  somewhat  in  water.     If  air  which  *s 
in  solution  in  water   is   pumped  out  and  analyzed,  it  is 
found  to  have  a  different  composition  from  that  of  or- 
dinary air.     Instead  of  containing  nearly  4  volumes  of 
nitrogen  to  1  of  oxygen,  it  contains  only  1.87  volumes  of 
nitrogen  to  1  of  oxygen.     The  proportion  of  oxygen  is 


252  INORGANIC  CHEMISTRY. 

much  larger  in  the  air  which  has  been  dissolved  in  water 
than  it  is  in  ordinary  air.  This  is  due  to  the  fact  that 
oxygen  is  more  soluble  in  water  than  nitrogen.  There- 
fore, when  air  is  shaken  with  water,  relatively  more  oxy- 
gen than  nitrogen  is  dissolved.  If  the  gases  were  in 
chemical  combination  we  should  expect  the  compound  to 
dissolve  as  such  and  without  change  of  composition. 

The  above  evidence  shows  that  nitrogen  and  oxygen 
are  not  combined  chemically  in  the  air,  but  that  they  are 
simply  mixed  together. 

As  enormous  quantities  of  oxygen  are  constantly  em- 
ployed in  the  processes  of  respiration  of  animals,  com- 
bustion, and  various  kinds  of  decay,  the  question  will 
suggest  itself  :  Is  the  quantity  of  oxygen  in  the  air  decreas- 
ing ?  In  regard  to  this  point  some  ingenious  calculations 
have  been  made  the  results  of  which  are  reassuring.  An 
approximate  estimate  of  the  extent  of  the  atmosphere, 
and  therefore  of  the  supply  of  oxygen,  can  easily  be  made. 
Assuming  that  the  population  of  the  earth  is  1000  million 
human  beings,  the  quantity  of  oxygen  used  by  them  in 
respiration  in  a  year  would  amount  only  to  about 
•swornnr  Par*  °f  ^ne  supply.  Suppose,  further,  that  for 
all  other  purposes  nine  times  as  much  oxygen  is  required 
as  for  the  respiration  of  human  beings,  then  the  total 
amount  used  up  in  a  year  would  be  only  g-g-gVoT  °^  *ne 
whole  supply.  In  3800  years  the  decrease  in  the  amount 
of  oxygen  in  the  air  would  be  only  1  per  cent.  Whether 
there  has  been  such  a  decrease  or  not  it  is  impossible  to 
say,  as  it  is  only  within  a  comparatively  few  years  that 
accurate  analyses  have  been  made.  It  appears  probable, 
however,  from  other  considerations  that  the  quantity  of 
oxygen  in  the  air  is  not  decreasing.  It  is  known  that 
the  process  of  plant  life  involves  a  giving  off  of  oxygen 
which  comes  from  other  compounds.  The  plants  have 
the  power  to  decompose  the  carbon  dioxide  found  in  the 
air,  and  they  utilize  the  carbon  and  a  part  of  the  oxygen, 
but  another  part  they  give  back  to  the  air,  so  that  in  the 
process  of  vegetable  growth  we  have  a  constant  source 
of  supply  of  oxygen. 


ANALYSIS  OF  AIR.  253 

Analysis  of  Air. — The  earliest  examinations  of  Priest- 
ley, Lavoisier,  and  Scheele  were  made  by  burning  sub- 
stances in  air  contained  in  closed  vessels.  They  con- 
cluded that  the  air  is  made  up  of  I  oxygen  and  f  nitrogen 
by  bulk.  In  order  to  determine  the  composition  of  the 
air  to-day,  we  should  proceed  as  follows  :  A  qualitative 
examination  would  easily  show  the  presence  of  nitrogen 
and  oxygen.  If  a  solution  of  calcium  hydroxide,  Ca(OH)2, 
which  is  known  as  lime-water,  or  a  solution  of  barium 
hydroxide,  Ba(OH)2,  is  exposed  to  the  air  it  becomes 
turbid,  and  a  precipitate  is  formed.  Neither  nitrogen  nor 
oxygen  nor  an  artificially  prepared  mixture  of  the  two 
gases  can  produce  this  change.  It  has  been  shown  that 
the  change  is  due  to  the  presence  in  the  air  of  a  small 
quantity  of  the  gaseous  compound,  carbon  dioxide,  CO2. 
If  calcium  chloride  or  phosphorus  pentoxide  is  exposed 
to  the  air  it  soon  becomes  moist  and  after  a  time  turns 
liquid.  This  effect  has  been  found  to  be  due  to  the  pres- 
ence of  water  vapor  in  the  air.  Nitrogen  obtained  from 
the  air  by  passing  the  latter  over  copper,  which  abstracts 
the  oxygen,  has  been  shown  to  contain  a  small  amount  of 
an  extremely  inert  gas,  argon  (which  see).  By  other 
methods  which  need  Hot  be  considered  here  it  can  be 
shown  that  there  are  many  other  substances  in  the  air 
besides  those  mentioned.  Among  them  are  ammonia, 
hydrogen  dioxide,  and  organic  matters  of  various  kinds, 
including  a  large  variety  of  germs  the  presence  of 
which  can  be  detected  by  the  changes  which  exposure  to 
the  air  produces  in  certain  liquids,  as  milk  and  fruit 
juices. 

Having  thus  learned  what  the  chief  constituents  of  the 
air  are,  the  next  thing  is  to  determine  in  what  quantities 
they  are  present,  or  to  make  a  quantitative  analysis  of 
the  air.  For  this  purpose  advantage  may  be  taken  of 
the  fact  that  phosphorus  when  exposed  to  the  air  at 
ordinary  temperatures  combines  slowly  with  the  oxygen, 
leaving  the  nitrogen.  If,  therefore,  a  piece  of  ordinary 
phosphorus  is  inserted  into  a  measured  volume  of  air 
contained  in  a  graduated  glass  tube  over  water  or  mer- 
cury, a  diminution  in  volume  will  take  place  slowly.  If, 


254  INORGANIC  CHEMISTRY. 

in  the  course  of  a  few  hours,  the  volume  is  again  meas- 
ured, the  difference  will  give  the  volume  of  oxygen 
absorbed,  while  the  gas  remaining  is  nitrogen.  Of 
course,  in  this  case  as  in  all  others  in  which  gas  volumes 
are  measured,  corrections  for  temperature,  pressure,  and 
tension  of  aqueous  vapor  must  be  made. 

Another  method  by  which  the  ratio  between  the  nitro- 
gen and  oxygen  in  air  can  be  determined  is  that  which 
was  first  employed  by  Dumas  and  Boussingault.  It  con- 
sists in  passing  air  over  heated  copper,  collecting  and 
measuring  the  nitrogen,  and  weighing  the  copper  oxide. 
The  apparatus  is  arranged  as  shown  in  Fig.  9. 


FIG.  9. 

The  copper  is  contained  in  the  glass  tube  ab  on  the 
combustion  furnace.  At  the  ends  of  this  tube  are  the 
stop-cocks  rr.  V  is  a  glass  globe  provided  with  a  stop- 
cock u.  Before  the  experiment  the  air  is  exhausted  from 
the  globe  and  the  tube  ab,  and  the  tube  then  carefully 
weighed.  The  tubes  B  and  C  and  the  apparatus  A  con- 
tain substances  which  have  the  power  to  absorb  the  car- 
bon dioxide  of  the  air.  The  tube  ab  is  now  heated  and 
air  admitted  after  passing  through  (7,  B,  and  A.  The 
copper  takes  up  the  oxygen,  and  the  nitrogen  enters  the 
globe  V.  After  the  globe  is  full  it  is  weighed,  then  ex- 
hausted and  weighed  again,  and  the  difference  gives  the 
weight  of  the  nitrogen.  The  tube  is  also  exhausted  and 
weighed,  and  the  difference  between  this  weight  and  that 


ANALYSIS  OF  AIR.  255 

of  the  exhausted  tube  before  the  experiment  gives  the 
weight  of  the  oxygen. 

The  most  refined  method  for  the  analysis  of  the  air  is 
the  eudiometric  method  of  Bunsen.  This  consists  in 
adding  some  pure  hydrogen  to  a  measured  volume  of  air 
contained  in  a  eudiometer  over  mercury,  and  then  ex- 
ploding the  mixture  by  means  of  an  electric  spark.  If 
the  conditions  are  right  all  the  oxygen  present  will  com- 
bine with  hydrogen,  and  in  consequence  of  this  there 
will  be  a  corresponding  contraction  in  the  volume  of  the 
gases.  The  amount  of  contraction  will  be  equal  to  the  vol- 
ume of  hydrogen  and  that  of  oxygen  which  have  combined 
to  form  water.  But  we  know  from  previous  experiments 
on  these  two  gases  that  they  combine  in  the  ratio  of  one 
volume  of  oxygen  to  two  of  hydrogen.  Consequently 
the  volume  of  oxygen  which  was  present  is  equal  to  one 
third  of  the  total  contraction.  Of  course  it  is  necessary 
that  there  should  be  enough  hydrogen  present  to  com- 
bine with  all  the  oxygen.  This  method  is  capable  of 
great  exactness.  The  most  accurate  analyses  made  by 
this  method  by  Bunsen  and  others  have  shown  that  in 
100  volumes  of  air  there  are  20.9  to  21  volumes  of  oxygen. 

The  estimation  of  the  quantity  of  water  vapor  present 
in  the  air  is  an  important  problem.  The  quantity  pres- 
ent depends  upon  a  variety  of  causes,  the  temperature 
and  the  direction  of  the  wind  being  the  chief  ones.  A 
good  chemical  method  for  estimating  the  water  consists 
in  drawing  a  known  volume  of  air  over  calcium  chloride 
in  a  weighed  tube.  This  substance  has  the  power  to 
take  up  water,  as  we  have  repeatedly  seen.  If  the  tube 
is  weighed  after  a  certain  volume  of  air  has  been  drawn 
through  it,  the  increase  in  weight  will  show  the  weight 
of  water  contained  in  that  volume  of  air. 

The  quantity  of  water  vapor  present  in  the  air  varies  be- 
tween comparatively  wide  limits.  At  any  given  temper- 
ature the  air  cannot  hold  more  than  a  certain  quantity. 
When  it  contains  this  quantity  it  is  said  to  be  saturated. 
If  cooled  down  below  this  temperature  the  vapor  partly 
condenses,  and  appears  now  as  water.  When  a  vessel 
containing  ice  is  placed  in  the  air,  that  which  immedi- 


256  INORGANIC  CHEMISTRY. 

ately  surrounds  the  vessel  is  cooled  down  below  the  point 
at  which  the  quantity  of  water  vapor  present  would  satu- 
rate the  air,  and  water  condenses  on  the  outside  of  the 
vessel.  Every  one  has  noticed  that  on  a  warm  cloudy  day 
more  water  condenses  on  such  a  vessel  than  on  a  clear 
cool  day.  The  water  vapor  present  in  the  air  has  an 
important  effect  on  man.  The  inhabitants  of  countries 
with  moist  climates  apparently  have  characteristics  which 
are  not  generally  met  with  in  those  who  inhabit  countries 
with  dry  climates.  The  difference  in  the  effects  of  moist 
and  of  dry  air  on  an  individual  is  well  known. 

When  air  which  is  charged  with  water  vapor  comes  in 
contact  with  cooler  air,  the  vapor  condenses  and  falls  as 
rain. 

The  method  employed  for  the  purpose  of  estimating 
the  quantity  of  carbon  dioxide  in  the  air  consists  in  draw- 
ing a  known  volume  of  air  over  something  which  has  the 
power  to  absorb  the  carbon  dioxide,  and  then  determin- 
ing the  increase  in  weight  of  the  absorbing  substance. 
Potassium  or  sodium  hydroxide  is  well  adapted  to  this. 
An  apparatus  has  been  constructed  in  which  barium 
hydroxide,  Ba(OH)3,  is  used  as  the  absorbent.  When 
carbon  dioxide  is  passed  through  a  solution  of  this  sub- 
stance insoluble  barium  carbonate,  BaCO3,  is  thrown 
down  according  to  the  equation 

Ba(OH)2  +  C02  =  BaC03  +  H2O. 

This  may  be  filtered  off  and  weighed,  and  the  quantity 
of  carbon  dioxide  estimated  from  the  results  ;  or,  if  a 
known  quantity  of  the  hydroxide  is  taken,  the  quantity 
left  unacted  upon  after  the  experiment  can  be  determined 
by  neutralizing  with  an  acid,  the  neutralizing  power  of 
which  has  previously  been  determined  with  care.  The 
quantity  of  carbon  dioxide  present  in  the  air  is  relatively 
very  small,  being  about  3  parts  in  10,000.  It  varies 
slightly  according  to  the  locality  and  season,  being 
greater  in  cities  and  in  summer  than  in  the  country  and 
in  winter ;  and  greater  in  warm  countries  than  in  cold. 
It  is  as  essential  to  the  life  of  plants  as  oxygen  is  to  the 
life  of  animals. 


AIR  AND  LIFE.  257 

It  is  not  an  easy  matter  to  determine  the  quantities  of 
the  other  constituents  of  the  air,  as  the  ammonia,  organic 
substances,  etc.,  though  there  is  no  difficulty  in  deter- 
mining that  they  are  present  in  very  small  quantities. 

The  relations  of  the  air  to  the  most  important  chemical 
changes  which  are  taking  place  upon  the  earth  form  one 
of  the  most  interesting  subjects  for  all  men.  We  have 
had  a  slight  glimpse  of  the  action  of  oxygen,  and  of  that 
of  carbon  dioxide  ;  both  are  essential  to  the  life  of  plants 
and  animals.  So,  too,  the  water  vapor  acts  chemically 
upon  plants,  and  probably  to  some  extent  in  the  respira- 
tion of  animals.  As  regards  the  nitrogen,  this  element  is 
frequently  referred  to  as  inert,  and  as  serving  the  purpose 
of  diluting  the  oxygen.  Inert  it  undoubtedly  is,  and 
there  is  also  no  doubt  that  it  dilutes  the  oxygen,  but 
these  statements  give  a  very  inadequate  conception  of 
the  important  part  played  by  it  in  the  processes  of  nature. 
Nitrogen  in  some  form  of  cpmbination  is  an  essential 
constituent  of  plants  and  animals.  The  animals  get  their 
nitrogenous  compounds  from  the  plants,  and  the  plants 
get  theirs  partly  at  least  from  the  soil.  By  the  growth 
of  plants,  therefore,  nitrogenous  compounds  are  con- 
stantly being  withdVawn  from  the  soil.  When  plants 
and  animals  undergo  decomposition  in  the  soil,  the  nitro- 
gen contained  in  them  is  gradually  converted  into  salts 
of  nitric  acid  or  nitrates,  and  if  the  decomposition  takes 
place  in  the  air  the  nitrogen  is  converted  principally  into 
ammonia.  Both  in  the  form  of  nitrates  and  of  ammonia 
the  nitrogen  can  be  utilized  by  plants,  so  that  if  the 
plants  and  animals  which  have  received  their  nourish- 
ment from  a  certain  tract  of  land  should  be  allowed  to 
decay  upon  this  land  after  death,  and  the  products  thus 
formed  should  be  uniformly  distributed  in  the  soil,  the 
latter  would  not  become  exhausted.  But  the  products 
of  the  soil  are  removed,  and,  therefore,  the  nitrogen  re- 
quired for  the  growth  of  other  plants  is  removed,  and 
the  soil  becomes  unproductive.  In  order  that  it  may  be 
rendered  fertile  again,  the  lost  nitrogen  must  be  sup- 
plied. It  appears  from  recent  very  elaborate  experi- 
ments that  the  plants  have  the  power  to  take  up  from 


258  INORGANIC  CHEMISTRY. 

the  air  a  part  of  the  nitrogen  which  they  need.  Whether 
they  take  it  up  directly,  or  it  is  first  taken  up  by  the 
soil  and  converted  into  some  compound  which  the  plants 
can  utilize,  are  questions  which  apparently  have  not 
been  answered  satisfactorily  as  yet.  It  appears  possible, 
further,  that  in  this  absorption  of  nitrogen  from  the 
air  certain  minute  organisms  which  exist  in  the  soil  play 
a  part. 

Pure  air  may  be  defined  as  air  which  consists  of  nitro- 
gen, oxygen,  and  carbon  dioxide  in  the  proportions  stated 
above,  together  with  some  water  vapor,  ammonia,  hy- 
drogen dioxide,  and  argon,  and  nothing  of  an  injurious 
nature.  It  is  evident  from  what  has  been  said  that  there 
is  constant  danger  of  contamination  from  natural  causes. 
The  most  common  cause  of  contamination  is  the  breathing 
of  human  beings  in  rooms  which  are  inadequately  sup- 
plied with  air.  The  breathing  process  involves  the  using 
up  of  oxygen  and  the  giving  off  of  carbon  dioxide  and 
small  quantities  of  organic  matter  which  is  undergoing  de- 
composition. If  the  quantity  of  oxygen  is  reduced  below 
a  certain  limit  the  air  becomes  unfit  for  breathing  pur- 
poses, and  evil  effects  follow.  An  ordinary  inhalation 
will  not  then  be  sufficient  to  supply  the  blood  with  the 
oxygen  necessary  to  purify  it,  and  the  system  will  begin 
to  suffer.  Headache,  drowsiness,  and  a  general  sense  of 
discomfort  follow.  The  ill  effects  of  breathing  the  air  of 
a  badly  ventilated  room  occupied  by  a  number  of  human 
beings  appear  to  be  due  for  the  most  part  to  the  pres- 
ence of  the  small  quantities  of  decomposing  organic  mat- 
ters which  are  given  off  from  the  lungs  with  the  carbon 
dioxide  and  other  gases.  These  act  as  poisons.  They 
have  been  thrown  off  from  the  lungs  because  they  are 
unfit  for  use,  and  when  they  are  taken  back  again  the 
normal  processes  of  the  body  are  interfered  with.  The 
subject  of  ventilation  has  been  so  thoroughly  discussed 
of  late  years  that  great  improvement  has  been  made 
in  the  arrangements  for  supplying  pure  air  to  dwelling 
apartments  and  audience  halls,  but  there  is  still  room  for 
improvement.  Fortunately,  in  most  buildings  there  is 
one  source  of  supply  of  pure  air  which  is  independent  of 


IMPURITIES  IN  AIR.  259 

architects'  plans.  This  is  the  diffusion  of  gases  through 
the  porous  materials  of  which  the  buildings  are  con- 
structed. This  diffusion  was  referred  to  under  the  head 
of  Hydrogen  (see  p.  45).  There  is  also  a  good  deal  of 
ventilation  through  the  cracks  and  other  apertures  which 
are  always  to  be  found  in  our  buildings. 

Under  some  conditions  not  thoroughly  understood  the 
air  becomes  contaminated  by  the  decomposition  of  ani- 
mal and  vegetable  matter.  Air  thus  contaminated  may 
cause  specific  diseases,  and  some  of  these  are  spoken  of 
as  being  caused  by  malaria,  a  word  which  signifies  sim- 
ply bad  air.  It  has  been  shown  that  there  are  in  the  air 
microscopic  germs  which  have  the  power  to  develop  in 
the  body,  and  then  to  cause  the  symptoms  which  are 
referred  to  malaria.  Disease  germs  and  germs  of  other 
kinds  are  present  in  the  air  in  great  variety,  and  they 
play  an  important  part  in  connection  with  the  life  and 
health  of  mankind. 

Argon. — Lord  Eayleigh  has  shown  that  a  liter  of  nitro- 
gen prepared  from  the  air  by  abstracting  the  oxygen  and 
the  small  quantities  of  other  substances  known  to  be 
present  weighs  1.2572  grams,  while  a  litre  of  nitrogen 
made  from  some  chemical  compound,  such  as  ammonia, 
weighs  1.2505  grams.  Chemically  prepared  nitrogen  is 
lighter  than  that  obtained  from  the  air.  This  observa- 
tion led  Lord  Rayleigh  and  "W.  Eamsay  to  a  more 
thorough  chemical  examination  of  the  air,  the  result  of 
which  was  the  discovery  of  a  constituent  previously  un- 
known. This  is  the  gas  argon.  It  can  be  obtained  by 
either  of  two  methods :  (1)  By  passing  air  over  heated 
copper  until  all  the  oxygen  is  abstracted,  and  then  over 
heated  magnesium  which  unites  with  the  nitrogen  but 
leaves  the  argon ;  (2)  By  mixing  air  with  oxygen  and 
passing  electric  sparks  through  the  mixture.  The  oxygen 
and  nitrogen  combine  and  the  product  can  easily  be  re- 
moved. After  all  the  nitrogen  has  thus  been  removed 
argon  remains  behind.  Argon  is  an  element  with  the 
atomic  weight  about  40.  It  is  present  in  the  air  to  the 
extent  of  about  1  per  cent  of  the  nitrogen.  It  cannot  be 
made  to  combine  with  any  other  element. 


CHAPTER  XVI. 

COMPOUNDS   OF  NITROGEN  WITH  HYDROGEN— WITH 
HYDROGEN  AND  OXYGEN— WITH  OXYGEN,  ETC. 

General  Conditions  which  give  Rise  to  the  Formation  of 
the  Simpler  Compounds  of  Nitrogen. — We  have  seen  that 
nitrogen  is  an  inactive  element,  showing  little  tendency 
to  combine  with  other  elements.  It  is  nevertheless  an 
easy  matter  to  get  compounds  of  nitrogen  with  many 
other  elements,  and  among  these  compounds,  some  of 
those  which  it  forms  with  hydrogen  and  oxygen  are  of 
much  importance. 

When  a  compound  which  contains  carbon,  hydrogen,  and 
nitrogen,  and  is  not  volatile,  is  heated  in  a  closed  vessel, 
so  that  the  air  does  not  have  access  to  it,  the  nitrogen 
passes  out  of  the  compound,  not  as  nitrogen,  but  partly 
in  combination  with  hydrogen,  in  the  form  of  the  com- 
pound ammonia.  Nearly  all  animal  substances  contain 
carbon,  hydrogen,  oxygen,  and  nitrogen  in  many  forms 
of  combination,  some  of  which  are  quite  complicated. 
Many  of  these  give  off  ammonia  when  heated.  Similarly, 
compounds  containing  carbon,  oxygen,  and  hydrogen, 
even  though  they  are  thoroughly  dry,  when  heated  give 
off  oxygen  in  combination  with  hydrogen  in  the  form  of 
water.  Both  these  kinds  of  decomposition,  that  which 
gives  ammonia  and  that  which  gives  water,  are  to  be 
ascribed  to  the  fact  that  the  compounds  of  carbon  which 
are  heated  are  unstable  at  higher  temperatures,  and 
when  they  are  broken  down  the  elements  contained 
in  them  arrange  themselves  in  combination  in  stable 
forms,  such  as  the  comparatively  simple  compounds, 
water  and  ammonia.  An  illustration  of  this  kind  of 
action  was  referred  to  in  speaking  of  the  preparation  of 
nitrogen  by  heating  ammonium  nitrite.  This  compound 

(260) 


COMPOUNDS  OF  NITROGEN".  261 

breaks  down  very  easily  under  the  influence  of  heat,  the 
hydrogen  and  oxygen  combining  to  form  the  stable  com- 
pound water,  while  the  nitrogen  remains  uncombined. 
Some  animal  substances,  as,  for  example,  urine,  give  off 
ammonia  when  they  undergo  spontaneous  decomposition 
in  the  air.  This  decomposition  is  generally,  if  not  always, 
due  to  the  action  of  minute  organisms,  the  germs  of 
which  are  in  the  air,  which  develop  when  they  come  in 
contact  with  certain  substances.  The  coal  which  is  used 
for  making  illuminating  gas  contains  some  hydrogen  and 
nitrogen  in  chemical  combination,  and  when  the  coal  is 
heated  ammonia  is  given  off  with  the  other  products. 

When  animal  substances  undergo  decomposition  in 
the  presence  of  basic  compounds  where  the  temperature 
is  comparatively  high,  the  nitrogen  combines  with  oxygen 
and  with  the  metal  of  the  base.  Either  a  salt  of  nitric 
acid,  HNO3,  or  of  nitrous  acid,  HNO2,  is  formed.  In 
some  countries  where  the  conditions  are  favorable  to  the 
process,  immense  quantities  of  nitrates  are  found,  chiefly 
potassium  nitrate,  KNO3,  and  sodium  nitrate,  NaNO3. 
Nitrates  are,  however,  found  everywhere  in  the  soil.  The 
change  of  animal  and  vegetable  nitrogenous  substances 
to  the  form  of  nitrates  is  probably  caused  by  the  action 
of  minute  living  organisms,  which  are  found  everywhere, 
and  serve  an  important  purpose  in  converting  the  waste 
animal  and  vegetable  matter  into  simple  compounds 
which  can  be  utilized  by  plants.  How  they  effect  the 
change  is  not  known.  From  the  salts  of  nitric  acid 
which  are  found  in  nature,  nitric  acid  itself  can  easily  be 
made. 

Nearly  all  the  compounds  of  nitrogen  with  which  we 
have  to  deal  are  made  either  from  ammonia  or  from 
nitric  acid. 

Relations  between  the  Principal  Compounds  of  Ni- 
trogen.— In  studying  the  compounds  of  sulphur,  we  saw 
that  whenever  a  compound  of  sulphur  is  oxidized  with  a 
strong  oxidizing  agent  the  final  product  of  the  action  is 
sulphuric  acid ;  and  that,  on  the  other  hand,  under  the 
influence  of  strong  reducing  agents  these  compounds 
yield  hydrogen  sulphide  as  the  final  product.  So  also 


262  INORGANIC  CHEMISTRY. 

the  limit  of  oxidizing  action  in  the  case  of  chlorine  is 
perchloric  acid,  and  of  reducing  action,  hydrochloric 
acid.  The  other  compounds  of  sulphur  with  hydrogen 
and  oxygen  are  products  intermediate  between  the  two 
limits,  sulphuric  acid  and  hydrogen  sulphide,  as  the 
other  compounds  of  chlorine  with  hydrogen  and  oxygen 
are  intermediate  between  hydrochloric  acid  and  per- 
chloric acid.  The  limit  of  reduction  of  nitrogen  com- 
pounds is  ammonia,  NH3,  and  of  oxidation,  nitric  acid, 
HNO3.  As  has  been  noticed,  the  valence  of  nitrogen 
towards  oxygen  is  greater  than  it  is  towards  hydrogen, 
but  the  difference  is  not  as  marked  as  in  the  case  of  the 
members  of  the  chlorine  group,  and  in  that  of  the  sulphur 
group.  Its  hydrogen  valence  is  3,  its  maximum  oxygen 
valence  is  5.  For  reasons  similar  to  those  which  have 
already  been  discussed  under  the  head  of  Hydroxides, 
the  oxidation  products  of  ammonia  are  believed  to  have 
their  oxygen  in  combination  with  hydrogen  in  the  form 
of  hydroxyl,  so  far  as  these  two  elements  are  present 
in  the  right  proportions  to  form  this  group.  If  this 
is  true  the  first  product  of  oxidation  of  ammonia 
would  have  the  structure  represented  by  the  formula 

(OH 
NK  H    .  -  A  compound  of  the  composition  represented  by 

(  H 

NH3O  is  known,  and  it  is  believed  that  it  has  the  above 
structure.  The  next  product  of  oxidation  formed  by  the 

(OH 
same  process  would  have  the  formula  NK  OH.    No  such 

(H 

compound  is  known,  however.  If  a  compound  of  this 
structure  should  lose  water  the  product  would  have  the 
composition  NOH,  and  probably  the  structure  O=N-H. 
A  compound  of  the  composition  represented  by  this 
formula  is  known,  but  it  seems  probable  from  its  con- 
duct that  the  formula  should  be  doubled,  N2O2H2,  and 
that  in  the  compound  there  are  two  hydroxyl  groups,  as 

N-OH 
represented  in  the  formula  II          .     Continuing  the  pro- 

N-OH 

cess  of  oxidation  of  ammonia,  the  next  product  which 


RELATIONS  BETWEEN  COMPOUNDS  OF  NITROGEN.    263 

we  should  expect  would  be  the  trihydroxyl  derivative 

(OH 
NS  OH.    Salts  derived  from  an  acid  of  this  composition, 

(OH 

which  should  be  called  normal  nitrous  acid,  are  known. 
But  this  normal  acid  breaks  down  readily  by  loss  of  water 

into  an  acid  of  the  formula  N  \  QTT     or     O=N-O-H, 

which  is  the  acid  from  which  nearly  all  the  nitrites  are 
derived.  By  continued  oxidation  the  quinquivalence  of 
the  nitrogen  makes  it  possible  to  add  another  oxygen 
atom  to  the  molecule  of  nitrous  acid,  and  thus  to  form 
the  final  product  of  oxidation,  nitric  acid  of  the  structure 
O 

N-O-H ;  or,  if  the  oxidation  takes  place  in  solution,  it 
I! 

O 
is  probable  that  the  final  product  is  a  hydroxyl  deriva- 

roH 

I  OH 
tive  of  the  structure  N-{  OH,  which  should  be  called 

I  OH 

LOH 

normal  nitric  acid.  There  are  some  salts  known  which 
are  derived  from  an  acid  of  this  composition,  but  most 
of  the  nitrates  are  derived  from  an  acid  which  is  formed 
from  this  by  loss  of  two  molecules  of  water : 

N(OH)5  =  N02(OH)  +  2H2O. 

It  is  impossible  at  present  to  furnish  satisfactory  evi- 
dence in  favor  of  the  views  just  pointed  out.  All  that 
can  be  said  is  that  the  views  are  in  accordance  with  a 
number  of  facts,  and  that  they  simplify  the  study  of  the 
relations  of  the  compounds. 

When  nitric  acid  is  reduced,  of  course,  the  above  pro- 
cesses must  be  regarded  as  reversed.  Thus,  the  first 
result  of  the  action  of  nascent  hydrogen  upon  nitric  acid, 
for  example,  would  be  the  elimination  of  one  atom  of 
oxygen  and  the  formation  of  nitrous  acid  : 

NO2(OH)  +  2H  =  NO(OH)  +  HaO. 


264  INORGANIC  CHEMISTRY. 

The  next  result  would  be  the  formation  of  the  compound 
N(OH)  or  N2O2H2  ;  the  next  the  formation  of  the  com- 
pound NH2(OH)  ;  and  finally  this  would  be  converted 
into  ammonia.  But  the  matter  is  complicated  by  the 
fact  that  some  of  the  compounds  referred  to  above  break 
down  into  water  and  oxides  of  nitrogen.  Thus,  nitrous 
acid,  NO(OH),  breaks  down  into  water  and  nitrogen  tri- 
oxide  or  nitrous  anhydride,  N2O3  ;  and  the  compound 
N2O2H2,  or  hyponitrous  acid,  breaks  down  into  water  and 
nitrous  oxide,  N2O  : 


N(OH) 

II  =  N2O  +  H2O. 

N(OH) 

Further,  there  are  some  compounds  of  nitrogen  and 
oxygen  for  which  there  are  no  analogous  hydroxyl  com- 
pounds known,  as,  for  example,  NO  and  NO2  or  N2O4.  In 
these  compounds  nitrogen  appears  to  be  bivalent  and 
quadrivalent,  while  in  nitrous  oxide,  N2O,  it  appears  to 
be  univalent.  Only  those  oxygen  compounds  of  nitro- 
gen in  which  the  nitrogen  is  univalent,  trivalent,  or 
quinquivalent  appear  to  form  corresponding  hydroxyl 
compounds.  Taking  the  simplest  view  of  the  matter, 
nitrogen  appears  to  be  univalent,  bivalent,  trivalent, 
quadrivalent,  and  quinquivalent  towards  oxygen,  as 
shown  in  the  compounds 

N2O,      NO,      N2O3,      NO2,      and      N205. 

Corresponding  to  nitrous  oxide,  N2O,  nitrogen  trioxide, 
N2O3,  and  nitrogen  pentoxide,  N2O5,  there  are  hydroxyl 
derivatives  which  have  acid  properties  : 

N20  +  H20  =  N2(OH)2; 
N2O3  +  H20  =  2NO(OH)  ; 
N206  +  H20  =  2N02(OH). 

But  there  are  no  hydroxyl  derivatives  corresponding 
to  those  oxides  in  which  the  valence  appears  to  be  2  and 
4.  Considering  the  ease  with  which  those  hydrox- 


RELATIONS  BETWEEN  COMPO  UNDS  OF  NITROGEN.    265 

ides  of  nitrogen  break  down  which  contain  two  hydroxyl 
groups  in  the  molecule,  the  fact  that  hydroxides  of  the 
formulas  N(OH)2  and  N(OH)4  do  not  exist  is  not  surpris- 
ing. We  should  expect  them  to  break  down  spontane- 
ously into  NO  and  NOa. 

The  methods  by  which  it  is  possible  to  pass  from  one 
of  the  oxides  or  hydroxides  of  nitrogen  to  the  other  will 
be  pointed  out  below.  A  tabular  list  of  the  compounds 
is  first  given : 

Oxides.  Acids. 

Nitrous  oxide,  N3O  Hyponitrous  acid,  N2(OH), 

Nitric  oxide,  NO(N2O2)  Nitrous  acid,  NO(OH) 

Nitrogen  trioxide,       )  XT  r>       Nitric  acid,  NO2(OH) 
(Nitrous  anhydride),  f1^3 
Nitrogen  peroxide,  NO2(N2O4) 


Nitrogen  pentoxide,  )  v  r| 
(Nitric  anhydride),    j     2    * 


Besides  the  above  there  are  the  basic  compounds, 
hydroxylamine,  NH2(OH),  ammonia,  NH3,  and  hydrazine, 
N2H4 ;  and  with  water  ammonia  probably  forms  the 
hydroxide  NH4(OH),  which  is  a  base.  In  addition  to 
these  there  is  a  compound  of  nitrogen  with  hydrogen  of 
the  formula  NSH.  This  is  a  strong  acid.  (See  Triazoic 
Acid.)  And  of  the  same  composition  as  hyponitrous 
acid  but  differing  from  it  in  properties  is  nitramide, 
N2H2O2,  the  relation  of  which  to  hyponitrous  acid  is  not 
yet  quite  clear. 

With  the  members  of  the  chlorine  group  nitrogen 
forms  a  few  compounds  which  are  characterized  by 
marked  instability.  So  unstable  are  they  that  they  ex- 
plode violently  when  simply  touched.  The  chloride  of 
nitrogen  explodes  with  terrific  violence  when  the  direct 
rays  of  the  sun  are  allowed  to  shine  upon  it.  With  sul- 
phur, nitrogen  forms  two  compounds.  With  sulphur, 
hydrogen,  and  oxygen,  however,  nitrogen  forms  a  number 
of  compounds,  one  of  which  has  already  been  referred  to 
in  connection  with  the  manufacture  of  sulphuric  acid. 
This  is  the  so-called  nitrosyl-sulphuric  acid  which  is 
formed  by  the  action  of  a  mixture  of  the  peroxide,  NO,, 


266  INORGANIC  CHEMISTRY. 

and  the  monoxide,  NO,  on  sulphurous  acid,  oxygen,  and 
water. 

In  studying  the  compounds  of  nitrogen,  it  will  be  best 
to  begin  with  the  end  products,  ammonia  and  nitric  acid. 

Ammonia,  NH3.  —  The  conditions  under  which  ammonia 
is  formed  have  been  mentioned.  The  chief  source  at 
present  is  the  "  ammoniacal  liquor"  of  the  gas-works. 
This  is  the  water  through  which  the  gas  has  been  passed 
for  the  purpose  of  removing  the  ammonia,  and  it  contains 
ammonia,  or  ammonium  hydroxide,  NH4(OH),  in  solu- 
tion. By  adding  hydrochloric  acid  to  this  solution  the 
salt  ammonium  cJdoride,  NH4C1,  is  formed.  This  is  the 
well-known  substance  sal  ammoniac.  It  appears  that  this 
name  has  its  origin  in  the  fact  that  common  salt  or 
sodium  chloride,  NaCl,  was  formerly  called  sal  armenia- 
cum,  and  that  afterward,  through  a  misunderstanding, 
ammonium  chloride  came  to  be  known  by  the  same  name 
which  underwent  change  to  the  form  sal  ammoniacum,  or 
sal  ammoniac.  When  the  ammoniacal  liquor  is  treated 
with  sulphuric  acid,  ammonium  sulphate  is  formed. 
From  one  or  the  other  of  these  salts  it  is  a  simple  matter 
to  obtain  ammonia.  For  this  purpose  it  is  only  necessary 
to  treat  the  salt  with  some  strongly  basic  compound,  as, 
for  example,  potassium  or  sodium  hydroxide,  or  calcium 
hydroxide.  Thus,  when  a  solution  of  potassium  hydrox- 
ide is  poured  on  ammonium  chloride  or  sulphate  the 
strong  penetrating  odor  of  ammonia  is  at  once  noticed. 
The  first  reaction  probably  results  in  the  formation  of 
ammonium  hydroxide,  thus  : 


NH4C1       +KOH       =NH4OH    +  KC1  ; 
(NH4)2SO4  +  2KOH     =  2NH4OH  +  K2SO4  ; 
2NH4C1     +  Ca(OH)2  =  2NH4OH  +  CaCl2  ; 
(NH4)2SO4  +  Ca(OH)2  =  2NH4OH.+  CaSO4. 

But  the  ammonium  hydroxide  breaks  down  very  readily 
into  water  and  ammonia,  which  escapes  as  a  gas  : 


Further,   ammonia   can   also   be    made    by   bringing 


AMMONIA.  26? 

nascent  nitrogen  and  nascent  hydrogen  together,  as,  for 
example,  by  heating  a  mixture  of  iron  filings,  potassium 
nitrate,  and  potassium  hydroxide.  Under  these  circum- 
stances the  iron  sets  hydrogen  free  from  the  hydroxide, 
and  nitrogen  from  the  nitrate,  and  they  unite  to 
form  ammonia.  The  formation  of  ammonia  by  reduction 
of  nitric  acid  can  be  shown  by  treating  some  granulated 
zinc  with  dilute  sulphuric  acid,  and  while  the  action  is 
in  progress  adding  nitric  acid  drop  by  drop.  The  nitric 
acid  is  thus  reduced,  and  the  ammonia  which  is  formed 
remains  in  combination  with  the  sulphuric  acid  as  am- 
monium sulphate.  Other  interesting  modes  of  formation 
of  ammonia  are  by  the  action  of  electric  sparks  on  nitro- 
gen in  the  presence  of  water,  and,  in  general,  by  the 
evaporation  of  water : 

N2  +  2H2O  =  NH4N02. 

The  product  ammonium  nitrite  is  always  found  in  the 
air  in  small  quantities. 

In  the  laboratory  ammonia  is  prepared  by  treating 
ammonium  chloride  with  slaked  lime  or  calcium  hydrox- 
ide. The  two  are  mixed  in  the  proportion  of  two  parts 
of  slaked  lime  to  one  of  ammonium  chloride,  placed  in  a 
flask  and  gently  heated,  when  the  ammonia  is  given  off 
at  once. 

It  is  frequently  more  convenient  to  heat  a  strong 
aqueous  solution  of  ammonia,  such  as  is  found  in  every 
chemical  laboratory.  Such  a  solution  when  gently 
heated  readily  gives  off  ammonia. 

Ammonia  is  a  colorless,  transparent  gas  with  a  very 
penetrating,  characteristic  odor.  In  concentrated  form 
it  causes  suffocation.  Its  specific  gravity  is  0.589;  that 
is  to  say,  it  is  but  little  more  than  half  as  heavy  as  air. . 
A  liter  of  the  gas  under  standard  conditions  weighs 
0.7635  gram.  It  can  easily  be  compressed  to  the  liquid 
form  by  pressure  and  cold.  When  the  pressure  is  re- 
moved from  the  liquefied  ammonia  it  passes  back  to  the 
form  of  gas,  and  in  so  doing  it  absorbs  a  great  deal  of 
heat.  These  facts  are  taken  advantage  of  for  the  arti- 


268  INORGANIC  CHEMISTRY. 

ficial  preparation  of  ice.     This  application  will  be  clear 
from  the  following  explanation  and  Fig.  10. 

An  aqueous  solution  of  ammonia  saturated  at  0°  is 
brought  into  the  strong  iron  cylinder  A,  and  then  gently 
warmed,  while  the  vessel  B  is  cooled  by  cold  water. 
The  gas  given  off  from  A  passes  though  the  bent  tubes 
into  B,  where  it  is  condensed 
to  a  liquid.  The  cylinder 
A  is  now  placed  in  a  ves- 
sel of  cold  water,  and  the 
water  which  is  to  be  frozen  is 
placed  in  a  cylinder  D,  and 
this  into  the  hollow  space  E 
in  the  vessel  B.  The  liquid 
ammonia  passes  rapidly  into 
the  form  of  gas  which  is  ab- 
sorbed in  the  water  in  A, 
while  at  the  same  time  so  FIG.  10. 

much  heat  is  absorbed  that  the  water  in  D  is  frozen. 

Ammonia  does  not  burn  in  the  air,  but  does  burn  in 
oxygen  with  a  pale  yellowish  flame.  It  is  absorbed  by 
water  in  very  large  quantity.  One  volume  of  water  at 
the  ordinary  temperature  dissolves  about  600  volumes  of 
ammonia  gas,  and  at  0°  about  1000  volumes.  The  sub- 
stance with  which  we  commonly  have  to  deal  under  the 
name  of  ammonia  is  a  solution  of  ammonia  in  water.  It 
is  called  "  spirits  of  hartshorn"  in  common  language.  The 
solution  has  the  odor  of  the  gas.  It  loses  all  its  gas  when 
heated  to  the  boiling  temperature.  The  solution  shows 
a  strong  alkaline  reaction,  and  has  the  power  to  neutral- 
ize acids  and  form  salts.  The  conduct  of  the  solution  is, 
in  fact,  strikingly  like  that  of  sodium  and  potassium 
hydroxides,  and  it  is  believed  that  in  the  solution  there 
is  contained  a  compound  of  the  formula  NH4(OH),  known 
as  ammonium  hydroxide,  and  formed  by  the  direct  action 
of  ammonia  upon  water.  If  this  is  true,  then  the  action 
of  ammonia  upon  acids  is  to  be  explained  as  follows : 
Ammonium  hydroxide  is  analogous  to  potassium  hydrox- 
ide, but  differs  from  it  in  that  it  contains  the  group  of 
atoms  NH4  in  place  of  the  atom  K.  In  some  way  this 


AMMONIUM  SALTS.  269 

group  plays  in  the  salts  formed  by  ammonia  the  same 
part  that  the  elementary  atom  potassium  plays  in  the 
salts  of  potassium,  and  just  as  the  latter  are  called  potas- 
sium salts,  so  the  former  are  called  ammonium  salts. 
According  to  this  the  ammonium  salts  are  salts  which 
contain  the  group  NH4,  known  as  the  ammonium  group, 
in  place  of  the  hydrogen  of  the  acids.  They  are  formed 
by  direct  combination  of  ammonia  with  the  acids,  or  by 
the  action  of  ammonium  hydroxide  upon  acids.  The 
analogy  between  the  action  of  ammonium  hydroxide  and 
that  of  potassium  hydroxide  upon  acids  is  clearly  shown 
by  the  aid  of  the  following  equations  : 

K(OH)       +HC1     =KC1  +H2O; 

NH4(OH)  +HC1     =  NH4C1       +H2O; 
K(OH)       +  HN03=KN03        +H2O; 
NH4(OH)  +  HN03  =  NH4N03    +  H2O; 
2K(OH)     +  H2S04  =  K2S04        +  2H2O  ; 
2NH4(OH)  +  H2S04  =  (NH4)2S04  +  2H2O. 

The  formation  of  the  ammonium  salts  by  direct  action 
of  ammonia,  NH3,  upon  acids  is  represented  in  the  fol- 
lowing equations  : 


NH3    +HC1     = 

+  HNO3  =  NH4NO3  ; 


The  strong  tendency  of  ammonia  to  combine  directly 
with  acids  is  shown  by  bringing  two  uncovered  vessels, 
one  containing  a  solution  of  ammonia,  and  the  other  a 
solution  of  hydrochloric  acid,  near  each  other.  A  dense 
cloud  will  at  once  be  noticed,  if  the  solutions  are  concen- 
trated. This  is  due  to  the  direct  combination  of  the  gases 
which  escape  from  the  solutions. 

While  the  assumption  of  the  existence  of  the  group 
ammonium,  NH4,  in  the  ammonium  salts  is  of  great  ser- 
vice in  dealing  with  these  salts,  and  while  this  assump- 
tion appears  to  be  entirely  justified  by  the  facts,  no 
compound  of  this  composition  has  as  yet  been  isolated. 


270  INORGANIG  CHEMISTRY. 

The  name  ammonium  is  given  to  the  hypothetical  com- 
pound on  account  of  the  fact  that  it  evidently  plays  the 
part  of  a  metallic  element,  and  it  is  customary  to  give 
such  elements  names  ending  in  ium.  While,  further,  it 
is  generally  believed  that  the  compound  ammonium 
hydroxide,  NH4(OH),  is  formed  when  ammonia  dissolves 
in  water,  the  compound  itself  has  not  been  isolated, 
owing  to  its  instability  and  tendency  to  break  down  into 
ammonia  and  water.  On  the  other  hand,  some  very  in- 
teresting derivatives  of  this  hydroxide  have  been  isolated. 
There  is  one  of  these  which  is  derived  from  the  hydrox- 
ide by  the  replacement  of  the  four  hydrogen  atoms  of 
the  ammonium  by  groups  of  carbon  and  hydrogen 
atoms.  This  compound  is  stable  and  can  be  isolated  as 
a  hydroxide  of  the  general  formula  NE4(OH),  in  which 
E4  represents  the  groups  of  carbon  and  hydrogen  atoms. 
It  acts  almost  exactly  like  potassium  hydroxide. 

Composition  of  Ammonia.  —  By  oxidation  under  the 
proper  conditions  it  is  possible  to  convert  the  hydrogen 
of  ammonia  into  water  and  leave  the  nitrogen  in  the  free 
state.  As  water  and  nitrogen  are  the  only  products 
formed,  and  the  quantity  of  oxygen  used  up  in  the  oxida- 
tion is  equal  to  the  quantity  of  oxygen  found  in  the- 
water  formed,  it  follows  that  nitrogen  and  hydrogen  are 
the  only  elements  contained  in  ammonia. 

When  electric  sparks  are  passed  for  some  time  through 
a  mixture  of  nitrogen  and  hydrogen,  some  ammonia  is 
formed.  Conversely,  when  electric  sparks  are  passed  for 
a  time  through  ammonia,  nitrogen  and  hydrogen  are 
obtained. 

If,  in  the  oxidation  of  a  known  weight  of  ammonia,  the 
water  formed  and  the  nitrogen  left  uncombined  are  accu- 
rately determined,  it  will  be  found  that  in  ammonia  the 
elements  are  combined  very  nearly  in  the  proportion  of 
fourteen  parts  by  weight  of  nitrogen  to  three  parts  by  weight 
of  hydrogen.  Further,  the  molecular  weight  determined 
by  the  method  of  Avogadro  is  approximately  17.  There- 
fore, the  molecular  formula  of  ammonia  is  NH3,  the 
atomic  weight  of  nitrogen  being  14. 

The  proportion  by  volume  in  which  the  two  elements 


COMPOSITION  OF  AMMONIA.  271 

combine  can  be  determined  by  the  following  method  : 
A  glass  tube  closed  at  one  end  and  provided  with  a  glass 
stop-  cock  at  the  other  is  filled  with  pure  chlorine  gas. 
By  means  of  a  small  funnel  attached  to  the  open  end  a 
little  of  a  strong  aqueous  solution  of  ammonia  is  slowly 
introduced  into  the  tube.  Reaction  takes  place  at  once 
between  the  chlorine  and  the  ammonia,  according  to  the 
equation 


and  the  hydrochloric  acid  unites  with  ammonia  to  form 
ammonium  chloride  : 


The  entire  change  is  therefore  represented  by  the  equa- 
tion 

4NH,  +  301  =  N  +  3NH4C1. 


Hydrogen  and  chlorine  combine  in  equal  volumes,  as  we 
have  already  learned.  Now,  if  we  start  with  a  measured 
volume  of  chlorine,  and  add  ammonia  to  it  until  it  is  all 
used  up,  we  know  that  the  volume  of  hydrogen  which 
has  been  extracted  from  ammonia  is  equal  to  the  volume 
of  chlorine  with  which  we  started.  If  we  measure  the 
volume  of  nitrogen  left  over,  we  know  the  volume  of  the 
nitrogen  which  was  in  combination  with  a  volume  of 
hydrogen  equal  to  that  of  the  chlorine  originally  taken. 
This  experiment  has  been  tried  repeatedly,  and  it  has 
been  found  that  the  ratio  of  the  volume  of  nitrogen  to 
that  of  the  hydrogen  with  which  it  was  combined  is  as  1 
to  3.  The  tube  being  full  of  chlorine  at  the  beginning 
of  the  experiment,  it  will  be  found  to  be  one-third  full 
of  nitrogen  at  the  end.  Therefore,  in  ammonia  1  volume 
of  nitrogen  is  combined  ivith  3  volumes  of  hydrogen. 

The  experiment  just  referred  to  will  perhaps  be  better 


272 


INORGANIC  CHEMISTRY. 


understood  by  the  aid  of  the  accompanying  diagram. 
The  chlorine  in  the  tube  may  be  represented  as  made 
up  of  three  equal  parts  or  volumes.  Each  volume  of 
chlorine  combines  with  an  equal  volume  of  hydrogen, 
leaving  the  nitrogen  uncombined.  The  volume  of  nitro- 
gen left  is  only  one  third  of  that  of  the  chlorine,  or  for 
three  volumes  of  chlorine  there  is  one  volume  of  mitrogen : 


combine 
with 


leaving 


Therefore  in  ammonia  the  gases  nitrogen  and  hydrogen 
are  combined  in  the  proportion  of  1  volume  of  nitrogen 
to  3  volumes  of  hydrogen : 


Volume  relations  in  ammonia. 


Another  question  in  regard  to  the  volume  relations 
remains  to  be  answered,  and  that  is :  When  nitrogen  and 
hydrogen  unite  in  the  proportions  above  stated,  how 
many  volumes  of  ammonia  gas  do  the  four  volumes  of 


AMMONIUM  AMALGAM.  273 

the  constituents  form  ?  It  is  not  possible  to  determine 
this  by  direct  combination  of  the  two  gases,  but  am- 
monia can  be  decomposed  into  its  constituents  by  con- 
tinued passage  of  electric  sparks  through  it.  When  this 
is  done  it  is  found  that  after  the  decomposition  the 
gases  occupy  twice  the  volume  which  was  occupied  by 
the  ammonia.  It  appears,  therefore,  that  when  hydrogen 
and  nitrogen  combine  to  form  ammonia  the  volume  is 
reduced  to  one  half,  or,  what  is  the  same  thing,  when 
three  volumes  of  hydrogen  combine  with  one  volume  of 
nitrogen  the  four  volumes  form  two  volumes  of  ammonia 
gas. 

The  above  facts  have  already  been  commented  upon 
in  speaking  of  the  combination  of  gases  in  general ;  and 
it  has  been  shown  that  chlorine,  oxygen,  and  nitrogen 
combine  with  hydrogen  in  entirely  different  ways  : 

(1)  1  volume  of  chlorine   combines  with  1  volume  of 
hydrogen  to  form  2  volumes  of  hydrochloric  acid  gas. 

(2)  1  volume  of  oxygen  combines  with  2  volumes  of 
hydrogen  to  form  2  volumes  of  water  vapor. 

(3)  1  volume  of  nitrogen  combines  with  3  volumes  of 
hydrogen  to  form  2  volumes  of  ammonia  gas. 

What  the  cause  of  these  differences  is  we  do  not  know. 
In  some  way  these  facts  are  directly  connected  with  the 
law  of  Avogadro  that  equal  volumes  of  all  gases  contain 
the  same  number  of  molecules,  and  with  the  power  of 
the  atoms  of  chlorine,  oxygen,  and  nitrogen  to  combine 
with  one,  two,  and  three  atoms  of  hydrogen  respectively. 
When  one  atom  of  chlorine  unites  with  one  atom  of 
hydrogen  the  result  is  a  molecule.  So  also  when  one 
atom  of  oxygen  unites  with  two  atoms  of  hydrogen,  and 
when  one  atom  of  nitrogen  unites  with  three  atoms  of 
hydrogen,  the  result  in  each  case  is  a  molecule,  and,  ac- 
cording to  the  law  of  Avogadro,  a  gaseous  molecule 
whether  it  consists  of  one  atom  or  a  hundred  atoms  oc- 
cupies the  same  space. 

Ammonium  Amalgam. — A  very  curious  substance 
which  appears  to  consist  of  mercury  and  ammonium 
is  formed  when  a  solution  of  ammonium  chloride 
is  treated  with  a  compound  of  sodium  and  mercury 


274  INORGANIC  CHEMISTRY. 

known  as  sodium  amalgam.  The  action  is  thought  to 
take  place  thus  : 

2NH4C1  +  Na2Hg  =  2NaCl  +  (NH4)2Hg. 

The  product,  ammonium  amalgam,  is  unstable,  break- 
ing down  very  soon  into  ammonia,  hydrogen,  and  mer- 
cury. It  will  be  referred  to  again  and  somewhat  more 
fully  under  the  head  of  Mercury. 

Metallic  Derivatives  of  Ammonium  Compounds  and  of 
Ammonia.  —  Ammonia  acts  upon  a  number  of  metallic 
salts,  forming  with  them  complex  derivatives  which  in 
some  cases  appear  to  be  ammonia,  and  in  others  am- 
monium, in  which  one  or  more  hydrogen  atoms  are  re- 
placed by  metal.  Thus  mercuric  chloride,  HgCl2,  acts 
upon  ammonia,  forming  a  compound  of  the  formula 
HgClNH2.  This  appears  to  be  ammonium  chloride  in 
which  two  hydrogen  atoms  are  replaced  by  a  mercury 
atom,  or  as  mercuric  chloride  in  which  one  chlorine 
atom  is  replaced  by  the  group  NH2,  known  as  the  amide 
group.  According  to  the  latter  view  the  structure  of 

Cl 
this  compound  is  represented  by  the  formula  Hg<^-rr  . 

Similarly,  copper  chloride,  CuCl2,  forms  a  compound  the 
composition  of  which  is  represented  by  the  formula 
CuCl2.2NH3.  It  seems  probable  that  this  is  ammonium 
chloride  in  which  two  hydrogen  atoms  are  replaced  by 

an  atom  of  copper,  Cu<>jT3ni-     There  are  many  such 


compounds,  particularly  of  the  metals  mercury,  copper, 
cobalt,  and  platinum.  The  structural  formulas  above 
given  are  to  be  regarded  merely  as  suggestions.  Experi- 
mental evidence  in  favor  of  them  is  lacking. 

Structure  of  Ammonium  Compounds.  —  In  ammonia  ni- 
trogen is  unquestionably  trivalent.  But  what  takes  place 
when  ammonia  acts  upon  water,  upon  hydrochloric  acid, 
and  upon  acids  in  general?  What  structure  is  to  be 
ascribed  to  the  ammonium  compounds?  The  view 
generally  held  is  that  in  the  ammonium  compounds  ni- 
trogen is  quinquivalent.  According  to  this,  ammonia  is 
an  unsaturated  compound,  and  when  brought  in  contact 


STRUCTURE  OF  AMMONIUM  COMPOUNDS. 


275 


with  water,  hydrochloric  acid,  etc.,  it  saturates  itself. 
The  conception  is  represented  thus : 


H 


=  N^ 


HN03  = 


H 
H 
H  ; 
H 

01 

H 
H 

H   ; 
H 
OH 

H 
H 
H 
H 

[NO3 


In  all  the  resulting  compounds  the  nitrogen  is  believed 
to  be  quinquivalent.  It  must  be  confessed  that,  while 
this  is  a  convenient  hypothesis,  further  evidence  for  or 
against  it  is  desirable.  One  objection  which  may  be 
raised  to  it  is  this.  It  is  assumed  that  when  the  stable 
compound  hydrochloric  acid  acts  upon  ammonia  the 
hydrogen  separates  from  the  chlorine  and  both  combine 
with  nitrogen ;  but  for  nitrogen  alone  chlorine  has  very 
little  attraction.  This  objection  may  not  be  a  real  one. 
It  may  be  shown  that,  while  chlorine  has  for  nitrogen 
alone  very  little  attraction,  it  has  great  attraction  for 
nitrogen  which  is  in  combination  with  hydrogen. 

Hydrazine,  N2H4. — A  compound  closely  related  to  am- 
monia and  ammonium  has  recently  been  prepared  by  a 
complicated  method  which  cannot  be  explained  here. 
This  is  known  as  hydrazine.  Its  composition  and  mo- 
lecular weight  are  represented  by  the  formula  N2H4.  A 
large  number  of  derivatives  of  hydrazine  are  known,  and 
have  been  studied  exhaustively.  From  what  has  been 
learned  in  regard  to  them  it  appears  that  the  constitution 

of  hydrazine  is  represented  by  the  formula  J,**    Accord- 


276  INOEGANIC  CHEMISTRY. 

ing  to  this  it  bears  to  ammonia  a  relation  similar  to  that 
which  hydrogen  peroxide  is  believed  to  bear  to  water  : 

OH3  NH3 

OH  NH2 


Hydrazine  is  a  liquid  that  boils  at  113°.5  in  a  current 
of  hydrogen.  It  acts  upon  acids  much  as  ammonia  does, 
forming  the  hydrazine  salts. 

Hydroxylamine,  NH2(OH).  —  This  compound  is  pre- 
pared by  reducing  nitric  acid  : 

NO2(OH)  +  6H  =  NH2(OH)  +  2H2O. 

It  can  also  be  prepared  by  reduction  of  nitric  oxide  : 

2NO  +  6H  =  2NH2(OH). 

Hydroxylamine  is  a  solid  consisting  of  leaflets  or  hard 
needles.  It  melts  at  33°.05,  and  boils  at  58°  under  a 
pressure  of  22  mm.  When  the  water  solution  is  evap- 
orated, both  ammonia  and  hydroxylamine  pass  over  with 
the  water.  Its  salts  are  easily  obtained  by  treating  the 
solution  with  acids  : 

NH2(OH)  +  HC1     =  NH3(OH)C1  ; 

NH2(OH)  +  HNO3  =  NH.(QH)NQ,, 

From  the  method  of  formation  and  the  composition  it 
appears  that  these  salts  are  ammonium  salts  in  which 
one  hydrogen  of  the  ammonium  is  replaced  by  hydroxyl. 
They  should  therefore  be  called  Jiydroxyl-ammonium  salts. 
One  of  the  most  characteristic  properties  of  hydroxyl- 
amine is  the  ease  with  which  it  breaks  down  into  ammonia, 
nitrogen,  and  water  : 

3NH2OH  =  N3  +  NH3  +  3H2O. 

If  brought  in  contact  with  compounds  capable  of  reduc- 
tion, it  reduces  them,  the  nitrogen  in  these  cases  gener- 
ally combining  with  oxygen  to  form  nitrous  oxide,  N2O. 
The  reduction  of  cupric  oxide,  CuO,  takes  place  accord- 
ing to  the  equation 

2NH2(OH)  +  4CuO  =  N2O  +  2Cu20  +  3H20. 


NITRIC  ACID.  277 

By  nascent  hydrogen  hydroxylamine  is  reduced  to 
ammonia. 

Triazoic  Acid,  N3H. — This  compound,  which  is  also 
called  hydrazoic  and  hydronitric  acid,  can  be  made  by  a 
number  of  reactions  involving  the  use  of  complex  organic 
substances.  A  simpler  method  consists  in  passing  am- 
monia gas  over  heated  metallic  sodium,  and  then  passing 
nitrous  oxide,  NaO,  over  the  resulting  product,  which  is 
sodium  amide.  The  following  reaction  takes  place: 

NH2Na  +  N20  =  N3Na  +  H3O. 

By  dissolving  in  water  the  sodium  salt  thus  formed, 
treating  with  dilute  sulphuric  acid,  and  distilling,  a  solu- 
tion of  the  acid  is  obtained.  The  compound  is  a  color- 
less liquid  that  boils  at  37°.  It  has  a  fearfully  penetrat- 
ing odor,  and  produces  bad  effects  upon  one  who  inhales 
it.  It  is  a  strong  acid  resembling  hydrochloric  acid.  It, 
as  weljl  as  some  of  its  salts,  is  extremely  explosive.  Its 
ammonium  salt  formed  by  direct  union  with  ammonia  is 
interesting  as  it  has  the  composition  N,NH4  or  N4H4. 

Nitric  Acid,  HNO3. — This  important  chemical  compound 
was  first  made,  though  not  in  pure  condition,  about  the 
ninth  century  by  distilling  saltpeter,  copper  sulphate, 
and  alum.  The  name  nitric  acid  has  its  origin  in  the 
fact  that  the  compound  is  formed  from  niter.  In  German 
it  is  called  Salpetersdure,  which  literally  translated  is 
saltpeter  acid.  It  has  already  been  stated  that  the  salts 
of  this  acid,  particularly  the  potassium  and  sodium  salts, 
occur  very  widely  distributed  in  the  earth,  and  that  there 
is  a  great  accumulation  of  the  sodium  salt  in  South 
America,  whence  the  name  Chili  saltpeter.  Wherever 
organic  matter,  particularly  that  of  animal  origin,  under- 
goes spontaneous  decomposition  in  the  presence  of  basic 
substances,  nitrates  are  formed,  probably  in  consequence 
of  the  action  of  an  organism  known  as  the  nitrifying  fer- 
ment. This  process  of  nitrification  has  already  been  re- 
ferred to  in  a  general  way.  It  is  one  of  great  importance 
for  the  welfare  of  the  human  race,  and  indeed  of  most 


278  INORGANIC  CHEMISTRY. 

living  beings,  as  by  its  aid  the  useless  nitrogenous  sub- 
stances of  dead  plants  and  animals  are  converted  back 
into  the  useful  nitrates  which  in  the  soil  aid  the  pro- 
cesses of  plant  growth. 

Nitric  acid  can  be  formed  by  the  action  of  electric 
sparks  on  nitrogen  and  oxygen  in  the  presence  of  water. 
It  is  also  easily  formed  by  the  action  of  oxidizing  agents 
on  ammonium  compounds. 

Nitric  acid  is  always  prepared  by  treating  potassium 
or  sodium  nitrate  with  concentrated  sulphuric  acid.  When 
sodium  nitrate  is  treated  with  sulphuric  acid  action  takes 
place  thus : 

NaNO3  +  HaSO4  =  NaHSO4  +  HNO3. 

The  salt  formed  in  this  way  is  primary  sodium  sulphate,  or 
acid  sodium  sulphate.  If  sufficient  of  the  saltpeter  is 
present  and  the  temperature  is  raised,  a  second  reaction 
takes  place,  resulting  in  the  formation  of  normal  sodium 
sulphate : 

NaN03  +  NaHSO4  =  Na2SO4  +  HNO3. 

But  the  temperature  required  for  this  reaction  is  so  high 
that  a  considerable  part  of  the  nitric  acid  is  decomposed. 
In  the  preparation  of  nitric  acid,  therefore,  the  first  re- 
action is  the  one  used,  and  for  this  purpose  the  sub- 
stances are  brought  together  in  a  retort  in  the  proportion 
of  their  molecular  weights  (about  equal  weights),  and 
the  retort  gently  heated.  The  nitric  acid  distils  over 
slowly,  and  is  condensed  by  cooling  the  receiver. 

On  the  large  scale  the  acid  is  made  by  bringing  Chili 
saltpeter  and  concentrated  sulphuric  acid  together  in  cast- 
iron  cylinders  or  retorts. 

Nitric  acid  is  a  colorless  volatile  liquid.  It  begins  to 
boil  at  86°,  but  at  this  temperature  it  undergoes  partial 
decomposition  into  nitrogen  peroxide,  water,  and  oxygen  : 

2HNO3  =  2N02  +  H20  +  O. 

It  undergoes  the  same  change  slowly  when  exposed  to 
the  direct  rays  of  the  sun.  In  consequence  of  this  de- 


NITRIC  ACID.  279 

composition  the  distillate  collected  in  the  manufacture  of 
nitric  acid,  and,  in  general,  whenever  the  acid  is  distilled, 
always  contains  a  considerable  percentage  of  water,  and 
is  colored  more  or  less  yellow  by  the  nitrogen  peroxide 
present.  In  order  to  abstract  the  water  from  ordinary 
nitric  acid  it  is  mixed  with  concentrated  sulphuric  acid  and 
slowly  distilled ;  but  even  under  these  circumstances  the 
product  is  colored  in  consequence  of  some  decomposition, 
and  it  also  contains  some  water.  By  conducting  carbon 
dioxide  gas  through  the  gently  warmed  acid  the  nitrogen 
peroxide  can  be  removed,  and  in  this  way  an  acid  con- 
taining about  99.5  per  cent  of  the  compound  HNO3  has 
been  obtained. 

Pure  nitric  acid  is  a  very  active  substance  chemically. 
It  gives  up  its  oxygen  readily  and  is  itself  thus  reduced 
to  other  compounds  of  nitrogen  and  oxygen,  or  of  nitro- 
gen, oxygen,  and  hydrogen,  as  has  already  been  pointed 
out.  When  it  acts  upon  the  metals  it  forms  nitrates, 
metal  atoms  being  substituted  for  the  hydrogen.  Accord- 
ing to  the  conditions,  nitrogen  peroxide,  NO2,  nitrous 
acid,  HNO2,  nitric  oxide,  NO,  nitrous  oxide,  N2O,  nitro- 
gen, hydroxylamine,  NH2(OH),  and,  finally,  ammonia  are 
formed  by  reduction  of  the  acid.  The  formation  of  am- 
monia and  of  hydroxylamine  by  reduction  of  nitric  acid 
has  already  been  specially  referred  to.  Of  these  reac- 
tions, that  which  gives  nitric  oxide,  NO,  is  the  one  which 
commonly  takes  place  on  treating  metals  with  nitric 
acid.  The  oxides  NO2  and  N2O3  are  themselves  readily 
reduced  to  nitric  oxide. 

If  the  element  upon  which  the  acid  acts  has  not  the 
power  to  replace  the  hydrogen,  the  action  consists  in 
oxidation.  This  is  shown  in  the  action  of  strong  nitric 
acid  upon  tin,  phosphorus,  carbon,  sulphur,  etc.  In  each 
case  the  highest  oxidation-product  is  formed.  Tin  is 
converted  into  normal  stannic  acid,  Sn(OH)4 ;  phosphorus 
into  phosphoric  acid,  PO(OH)3 ;  carbon  into  carbon  di- 
oxide, COa ;  and  sulphur  into  sulphuric  acid,  SO2(OH)2. 
It  disintegrates  carbon  compounds  very  readily,  convert- 
ing them  into  their  final  products  of  oxidation.  In  con- 
tact with  the  skin  it  causes  bad  and  dangerous  wounds. 


280  INORGANIC  CHEMISTRY. 

Upon  some  stable  compounds  of  carbon  it  acts  forming 
so-called  nitro-compounds,  very  much  as  chlorine  acts 
upon  them  forming  chlorine  substitution-products,  as 
was  explained  under  Chlorine.  The  formation  of  a  nitro- 
compound  takes  place  as  represented  in  the  equation 

C.H.  +  HO.NO,  =  C9HS.NO,  +  H.O. 

The  compound  C6H6  is  benzene,  and  the  compound 
C6HB.NO2  is  nitro-benzene. 

The  acid  mostly  used  in  the  laboratory  has  the  specific 
gravity  1.2  and  contains  32  per  cent  nitric  acid,  HNO3. 
The  commercial  acid  contains  about  68  per  cent  of  the 
acid. 

When  a  mixture  of  nitric  acid  and  water  is  boiled 
under  the  ordinary  atmospheric  pressure  it  loses  either 
water  or  nitric  acid  until  it  contains  68  per  cent  of  the 
acid  which  passes  over.  This  does  not  correspond  to 
any  definite  hydrate  of  nitric  acid,  though  it  approxi- 
mates the  composition  required  by  normal  nitric  acid, 
N(OH)B,  or  HNO3  +  2H2O,  and  it  is  probable  that  this 
hydrate  is  the  chief  constituent  of  the  mixture. 

Nitric  acid  is  a  strong  monobasic  acid,  forming  salts 
of  the  general  formula  MNO3,  all  of  which  are  soluble 
in  water.  Because  nitric  acid  is  a  strong  acid,  and  all 
normal  nitrates  are  soluble  in  water,  nitric  acid  is  one  of 
the  best  solvents.  On  the  other  hand,  the  acid  forms 
basic  salts,  some  of  which  are  difficultly  soluble  or  in- 
soluble in  water.  An  example  of  such  insoluble  basic 


(N03 
U 


nitrates  is  the  nitrate  of  bismuth  of  the  formula  Bi  <  OH  , 

(OH 

which  is  to  be  regarded  as  bismuth  hydroxide  one-third 
neutralized  by  nitric  acid.  There  are  some  apparently 
complex  salts  of  nitric  acid  which  are  derived  from  the 
normal  acid,  N(OH)5,  as,  for  example,  the  salt  HPb3O6N, 


which  should  probably  be  expressed  thus,  N  -I  O    -p-,        , 

0 

[O.Pb.OH 

being  a  basic  lead  salt  of  the  normal  acid.     There  are, 


-RED  FUMING  NITRIC  ACID— NITROUS  ACID.        281 

further,  some  salts  which  are  derived  from  the  acid 
NO(OH)3,  which  is  formed  by  abstracting  one  molecule  of 
water  from  normal  nitric  acid  : 

N(OH)6  =  NO(OH)8  +  H,0. 

By  far  the  largest  number  of  nitrates,  however,  are '  de- 
rived from  the  acid  of  the  formula  HNO3. 

Ked  Fuming  Nitric  Acid  is  formed  in  the  manufacture 
of  nitric  acid  from  saltpeter  and  sulphuric  acid  if  the 
temperature  is  raised  to  a  sufficient  extent  to  cause  the 
acid  sulphate  to  act  upon  the  nitrate : 

NaN03  +  HNaSO4  =  Na2SO4  +  HNO3. 

At  this  temperature  the  nitric  acid  undergoes  consider- 
able decomposition.  The  nitrogen  peroxide  formed  is 
absorbed  by  the  nitric  acid,  and  the  product  thus  ob- 
tained is  the  red  fuming  acid.  It  acts  more  energetically 
than  nitric  acid,  and  finds  some  applications  in  the  lab- 
oratory and  in  the  arts.  When  heated  it  gives  off 
nitrogen  peroxide,  and  if  diluted  with  water  it  is  changed 
to  ordinary  nitric  acid,  as  nitrogen  peroxide  is  decom- 
posed by  water,  forming  nitric  acid  and  nitric  oxide  or 
nitrous  acid,  according  to  the  temperature  of  the  water. 

Nitro-Tiydrochloric  Acid  or  Aqua  Regia  is  a  liquid 
formed  by  mixing  concentrated  nitric  and  hydrochloric 
acids.  It  was  called  aqua  regia  because  it  can  dissolve 
gold,  the  king  of  the  metals.  The  active  power  of  this 
liquid  as  a  solvent  of  metallic  substances  is  due  to  the 
fact  that  it  gives  off  chlorine,  and  a  compound  of  nitrogen, 
oxygen,  and  chlorine  which  readily  gives  up  its  chlorine. 
This  compound  has  the  composition  represented  by 
the  formula  NOC1,  and  it  is  best  designated  by  the 
name  nitrosyl  chloride.  The  product  of  the  action,  of 
mtro-hydrochloric  acid  upon  a  metal  is  the  correspond- 
ing chloride. 

Nitrous  Acid,  HNO2. — When  certain  salts  of  nitric  acid 
are  reduced  they  yield  the  corresponding  nitrites.  Thus, 
when  potassium  nitrate  is  heated  with  metallic  lead 
this  reaction  takes  place : 

KNO3  +  Pb  =  KNOa  +  PbO. 


282  INORGANIC  CHEMISTRY. 

Indeed,  if  potassium  nitrate  is  heated  alone  it  loses 
oxygen  and  is  converted  into  the  nitrite  : 

2KN03  =  2KN02  +  O2 ; 

but  the  reaction  is  not  complete,  and  the  salt  thus  ob- 
tained always  contains  more  or  less  nitrate. 

If  an  attempt  is  made  to  isolate  nitrous  acid  from  a 
nitrite  the  product  is  the  anhydride,  nitrogen  trioxide, 
N2O3.  Thus,  if  sulphuric  acid  is  added  to  potassium 
nitrite  the  following  reaction  takes  place  : 

2KN02  +  H2S04  =  N203  +  H2O  +  K2SO4. 

It  may  be  that  the  first  action  is  the  liberation  of 
nitrous  acid,  and  that  this  then  breaks  down  by  loss  of 
water.  The  two  reactions  are  represented  thus : 

2KN02  +  H2S04  =  2HNO2  +  KaSO4 ; 
2HNO3  =  Na03  +  H20. 

A  certain  analogy  will  be  observed  between  this  action 
and  that  which  takes  place  in  the  action  of  potassium 
hydroxide  upon  an  ammonium  salt,  when  ammonium  hy- 
droxide is  probably  first  given  off,  and  then  breaks  down 
into  ammonia  and  water. 

Salts  are  known  which  are  derived  from  normal  nitrous 
acid,  N(OH)8,  but  most  of  the  nitrites  are  derived  from 
the  acid  of  the  formula  NO(OH),  which  is  to  be  regarded 
as  formed  from  the  normal  acid  by  loss  of  one  molecule 
of  water : 

N(OH)3  =  NO(OH)  +  H2O. 

Hyponitrous  Acid,  H2N2O2. — The  sodium  salt  of  this 
acid  is  made  by  reducing  sodium  nitrite  in  solution  by 
means  of  sodium  amalgam  : 

2NaN03  +  8H  =  Na2N,Oa  +  4HaO. 

The  acid  can  also  be  made  by  oxidation  of  hydroxyl- 
amine. 

It  is  a  solid  consisting  of  white  crystalline  plates.  It 
is  very  explosive  when  freed  from  water.  In  water  solu- 


NITROUS  OXIDE.  283 

tion  it  is  much  more  stable,  but  at  the  ordinary  tempera- 
ture it  breaks  down  gradually,  the  principal  products 
being  nitrous  oxide  and  water  : 


It  would  appear  from  this  that  nitrous  oxide,  N2O,  bears 
to  hyponitrous  acid  the  relation  of  an  anhydride,  but 
salts  of  the  acid  cannot  be  obtained  from  nitrous  oxide 
directly.  The  molecular  weight  of  hyponitrous  acid  in 
solution  has  been  determined,  and  the  result  shows 
that  the  acid  has  the  double  formula  H2NaO2.  Its 
structure  is  probably  to  be  represented  by  the  for- 

N(OH) 

mula  II 

N(OH) 

Nitrous  Oxide,  N2O.  —  This  compound  can  be  obtained 
by  reduction  of  nitric  acid,  and  is  sometimes  formed  in 
considerable  quantity  when  copper  is  treated  with  the 
concentrated  acid,  though  when  made  in  this  way  it 
is  always  mixed  with  a  large  proportion  of  nitric  oxide. 
The  best  way  to  make  it  is  to  heat  ammonium  nitrate, 
NH4NO3,  which  breaks  down  into  nitrous  oxide  and  water  : 

NH4N03  =  N20  -f  2H2O. 

In  the  same  way  we  have  seen  that  ammonium  nitrite 
breaks  down  into  free  nitrogen  and  water  when  heated  : 

NH4N02  =  N2  +  2H20. 

In  these  reactions  we  see  exhibited  the  tendency  of 
hydrogen  and  oxygen  to  combine  at  elevated  tempera- 
tures. At  ordinary  temperature  this  tendency  is  not 
strong  enough  to  cause  a  disturbance  of  the  equilibrium 
of  the  parts  of  the  compound.  As  the  temperature  is 
raised  and  the  equilibrium  thus  disturbed/the  affinity 
of  the  hydrogen  for  the  oxygen  asserts  itself.  The  two 
elements  combine  to  form  water,  and  the  decomposition 
above  represented  takes  place. 

Nitrous  oxide  is  a  colorless,  transparent  gas  which  has 


284  INORGANIC  CHEMISTRY. 

a  sweetish  taste  and  odor.  Its  specific  gravity  is  1.527. 
It  is  somewhat  soluble  in  water  ;  one  volume  of  water  at 
0°  dissolving  somewhat  more  than  its  own  volume  of  the 
gas.  It  supports  combustion  almost  as  well  as  pure 
oxygen.  Some  substances  which  burn  in  oxygen  do  not, 
however,  burn  in  nitrous  oxide.  Sulphur  which  burns 
in  oxygen  is  extinguished  in  nitrous  oxide,  unless  it  is 
previously  heated  to  a  high  temperature.  To  under- 
stand the  action  of  this  compound  in  supporting  com- 
bustion it  must  be  borne  in  mind  that,  when  anything 
burns  in  oxygen,  the  oxygen  molecules  must  first  be 
broken  down  into  atoms  before  the  combination  can  take 
place.  Thus,  when  carbon  and  oxygen  are  brought  to- 
gether we  have  at  first  a  condition  represented  by  these 
symbols  : 


the  question  as  to  the  condition  of  the  carbon  being  left 
open.  When  the  temperature  is  raised  to  a  sufficiently 
high  point  the  condition  is  represented  thus  : 

O  +  O  +  C; 
and  now  the  reaction  takes  place  : 

o  +  o  +  c  =  co2. 

In  the  act  of  burning  in  free  oxygen,  therefore,  there  is 
always  a  certain  resistance  to  be  overcome.  Now,  when  a 
combustible  substance  is  brought  into  a  gas  which  gives 
up  its  oxygen  easily  the  condition  is  much  like  that  in 
free  oxygen.  If  the  temperature  is  raised  the  gas  is  de- 
composed, and  the  oxidation  then  follows.  In  the  case 
of  nitrous  oxide  this  decomposition  takes  place  : 


and  this  oxygen  in  the  atomic  condition  effects  the  oxi- 
dation. 

When  inhaled,  nitrous  oxide  causes  a  kind  oi»  intoxica- 
tion, which  is  apt  to  show  itself  in  the  form  of  hysterical 
laughing.  Hence  the  gas  is  called  laughing  gas.  In- 


NITRIC  OXIDE.  285 

haled  in  larger  quantity  it  causes  unconsciousness 
and  insensibility  to  pain.  It  is  therefore  used  ex- 
tensively to  prevent  pain  in  some  surgical  operations, 
particularly  in  extracting  teeth. 

When  subjected  to  a  low  temperature  and  high  pres- 
sure the  gas  is  easily  liquefied,  and  enclosed  in  properly 
constructed  metallic  cylinders  the  liquid  is  now  sent  into 
the  market.  In  order  to  get  the  gas  it  is  only  necessary 
to  open  the  stop- cock  of  the  cylinder.  When  the  liquid 
comes  in  contact  with  the  air  it  rapidly  turns  to  gas,  and 
the  temperature  is  very  much  lowered  in  consequence. 
This  causes  a  part  of  the  liquid  to  solidify. 

Nitric  Oxide,  NO. — This  is  the  most  stable  compound 
of  nitrogen  and  oxygen,  and  is  the  common  product 
of  the  reduction  of  nitric  acid.  Thus,  when  nitric  acid 
acts  upon  copper  and  other  metallic  elements  the  chief 
product  is  generally  nitric  oxide,  though,  as  we  have 
seen,  the  reduction  may  be  carried  farther.  The  prin- 
cipal action  in  the  case  of  copper  is  represented  thus : 

8HNO,  +  3Cu  =  3Cu(N03)3  +  2NO  +  4H,O. 

Considering  the  ease  with  which  nitric  acid  gives  up 
its  oxygen,  and  the  ease  with  which  copper  takes  up 
oxygen,  it  is  probable  that  the  copper  abstracts  oxygen 
directly  from  the  acid  as  represented  thus : 

2HNO3  +  3Cu  =  CuO  +  H,O  +  2NO. 

In  this  case  the  copper  oxide  would  at  once  form  copper 
nitrate  with  the  excess  of  nitric  acid : 

6HNO3  +  3CuO  =  3Cu(NO3)2  +  3H2O. 

Or,  combining  the  two  equations,  the  total  action  is  rep- 
resented in  the  same  way  as  it  is  above.  The  nitric 
acid  must  not  have  a  specific  gravity  higher  than  1.2,  and 
the  temperature  must  be  kept  down,  otherwise  the  reduc- 
tion of  the  nitric  acid  is  carried  farther  and  considerable 
nitrous  oxide  is  formed. 


286  INORGANIC  CHEMISTRY. 

It  is  possible  that  to  some  extent  the  hydrogen  liber- 
ated from  the  acid  may  act  as  a  reducing  agent,  thus- 
causing  the  formation  of  the  lower  oxides  of  nitrogen, 
as,  for  example, 

2HN03  +  6H  =  2NO  +  4HaO. 

A  good  method  for  making  pure  nitric  oxide  consists  in 
treating  ferrous  chloride,  FeCl2,  with  saltpeter.  The  re- 
action is  represented  thus : 

3Fe012  +  KNO3  +  4HC1  =  3FeCl3  +  KC1  +  2H3O  +  NO. 

The  reducing  action  is  here  effected  by  the  ferrous 
chloride,  FeCla,  which  tends  to  pass  over  into  ferric 
chloride,  FeCl3,  in  the  presence  of  hydrochloric  acid,  and 
anything  which  has  the  power  to  take  up  hydrogen. 
With  hydrochloric  acid  alone  it  does  not  form  ferric 
chloride,  but  if  any  reducible  compound  is  present  ac- 
tion takes  place  thus : 

2FeCl2  +  2HC1  +  O  =  2FeCl3  +  H2O. 

In  the  above  reaction  saltpeter  furnishes  the  oxygen, 
and  it  is  consequently  reduced  and  breaks  down,  yielding 
nitric  oxide,  while  the  potassium  forms  potassium  chlo- 
ride with  chlorine. 

Nitric  oxide  is  a  colorless,  transparent  gas.  Its  most 
remarkable  property  is  its  power  to  combine  directly 
with  oxygen  when  the  two  are  brought  together.  The 
act  of  combination  is  not  accompanied  by  the  appear- 
ance of  light,  though  heat  is  evolved.  In  the  reaction 
which  takes  place  at  ordinary  temperatures  nitrogen 
peroxide,  NO, ,  is  formed : 

NO  +  O  =  NO2. 

The  product  is  a  colored  gas,  and  the  change  of  the  color- 
less nitric  oxide  to  this  colored  product  can  therefore 
easily  be  recognized.  This  reaction  is,  further,  the  chief 
cause  of  the  reddish-brown  fumes  seen  when  nitric  acid 
acts  upon  metals  and  other  elements.  At  a  low  temper- 
ature some  nitrogen  trioxide  is  formed  when  oxygen  acts 
upon  nitric  oxide. 


NITROGEN  TRIOXIDE.  287 

From  what  has  already  been  said,  it  will  appear  that 
in  nitric  oxide  the  oxygen  and  nitrogen  are  more  firmly 
united  than  in  the  other  oxides.  Most  burning  sub- 
stances  are  extinguished  when  introduced  into  it,  though 
a  few  when  heated  in  it  to  a  high  temperature  extract  all 
or  a  part  of  the  oxygen.  Zinc  and  iron  extract  half  the 
oxygen  and  convert  nitric  oxide  into  nitrous  oxide. 
Potassium  and  sodium  decompose  it,  leaving  the  nitrogen 
free. 

A  curious  reaction  by  means  of  which  it  is  possible  to 
separate  nitric  oxide  from  other  gases  takes  place,  when 
the  oxide  is  passed  into  a  solution  of  ferrous  sulphate, 
FeSO4.  Under  these  circumstances  an  unstable  dark- 
colored  compound  is  formed,  which  appears  to  have  the 
composition  FeSO4  -|-  2NO.  When  the  solution  contain- 
ing it  is  heated  the  pure  gas  is  given  off. 

By  nascent  hydrogen  nitric  oxide  is  reduced  to  am- 
monia and  hydroxylamine. 

Nitrogen  Trioxide,  N2O3.  —  This  oxide  is  formed  by  ad- 
dition of  oxygen  to  nitric  oxide  at  low  temperatures  ;  by 
decomposition  of  the  nitrites  by  means  of  acids  ;  and  by 
the  combination  of  nitric  oxide  with  the  peroxide  at  a 
temperature  below  —21°.  The  gas  given  off  when  nitric 
acid  is  reduced  with  starch  or  arsenious  oxide,  AsaO3  , 
appears  to  be  a  mixture  of  nitric  oxide  and  the  peroxide. 
Pure  nitrogen  trioxide  is  a  liquid  of  an  indigo-blue  color. 
At  a  temperature  below  0°  it  undergoes  partial  decom- 
position into"  nitrogen  peroxide  and  nitric  oxide  : 


With  cold  water  nitrogen  trioxide  undergoes  decomposi- 
tion accompanied  by  an  evolution  of  nitric  oxide.  Pos- 
sibly this  reaction  takes  place  : 

3N,O3  +HaO  =  2HN03  +  4NO. 

By  treating  the  oxide  with  a  solution  of  sodium  hydrox- 
ide or  potassium  hydroxide  the  corresponding  nitrite  is 
formed  : 

2KOH  +  N,O3  =  2KNO2  +  H,O. 


288  INORGANIC  CHEMISTRY. 

Nitrogen  Peroxide,  NO2.  —  When  nitric  oxide  and  oxygen 
are  brought  together  in  the  proportion  of  2  volumes  of 
the  former  to  1  volume  of  the  latter  they  combine  com- 
pletely to  form  nitrogen  peroxide.  These  relations  will 
be  readily  understood  when  it  is  borne  in  mind  that  2 
molecules  of  nitric  oxide  require  1  molecule  of  oxygen 
to  effect  the  change,  as  is  shown  in  the  equation 


The  compound  is  most  easily  obtained  by  heating  lead 
nitrate,  when  nitrogen  peroxide  and  oxygen  are  given  off, 
and  lead  oxide  remains  behind  in  the  vessel  : 

Pb(NO3)2  =  PbO  +  2NO2  +  O. 

If  the  gases  are  passed  through  a  tube  surrounded  by  a 
freezing  mixture  the  peroxide  is  condensed  to  the  form 
of  liquid,  while  the  oxygen  passes  on.  When  perfectly 
dry  the  peroxide  is  easily  solidified.  It  acts  energetic- 
ally upon  compounds  which  have  the  power  to  take  up 
oxygen.  When  treated  with  water  it  undergoes  decom- 
position. If  the  temperature  is  low,  nitrous  and  nitric 
acids  are  formed  : 

2N02  +  H2O  =  HNO2  +  HNO3. 

If  the  water  is  hot,  however,  the  products  are  nitric  acid 
and  nitric  oxide  : 

3NO2  +  H2O  =  2HNO3  +  NO. 

The  nitric  oxide  thus  formed  will  take  up  oxygen  from 
the  air  and  yield  nitrogen  peroxide  again,  and  this,  in 
contact  with  hot  water,  will  be  decomposed,  forming 
nitric  acid  and  nitric  oxide,  until  all  the  peroxide  is  con- 
verted into  nitric  acid. 

The  determinations  of  the  specific  gravity  of  the  gas 
from  the  peroxide  show  that  at  low  temperatures  the 
molecular  formula  is  NaO4,  but  that  when  the  tempera- 
ture 150°  is  reached  the  molecule  is  represented  by  the 
formula  NO2.  The  compound  appears  therefore  to 
undergo  gradual  decomposition  or  dissociation  by  heat, 
so  that  until  the  temperature  150°  is  reached  the  gas  is 
a  mixture  of  the  compounds  NaO4  and 


STRUCTURE  OF  COMPOUNDS  OF  NITROGEN.        289 

Nitrogen  Pentoxide,  N2O5.  —  This  compound,  which 
bears  to  nitric  acid  the  relation  of  an  anhydride,  is 
formed  by  passing  chlorine  over  silver  nitrate  and  con- 
densing the  product.  The  reaction  takes  place  thus  : 

2AgN03  +  01,  =  N,05  +  2AgCl  +  O. 

It  is  also  formed  by  treating  nitric  acid  with  phosphorus 
pentoxide,  P2O6,  a  compound  which  has  a  very  marked 
power  to  unite  with  water.  The  action  is  represented 
thus  : 

2HNO3  =  NA  +  H2O. 

The  pentoxide  is  a  crystallized  substance,  which  readily 
decomposes  into  nitrogen  peroxide  and  oxygen.  In  con- 
sequence of  the  ease  with  which  it  gives  up  its  oxygen  it 
acts  violently  upon  many  oxidizable  substances.  With 
water  it  forms  nitric  acid  : 

NA  +  H20  =  2HNO,. 

Structure  of  the  Compounds  of  Nitrogen  with  Oxygen 
and  Hydrogen.  —  Our  knowledge  of  the  structure  of  the 
compounds  with  which  we  have  just  been  dealing  is  un- 
satisfactory. There  is  at  present  no  way  of  deciding 
whether  in  a  compound  like  nitrous  oxide,  for  example, 
the  oxygen  is  in  combination  with  both  nitrogen  atoms  : 
there  are  no  reactions  of  the  compound  which  throw  any 
light  upon  this  question.  Similar  difficulties  are  met 
with  in  connection  with  the  other  compounds  of  nitrogen 
and  oxygen.  As  was  remarked  on  page  264,  the  simplest 
view  which  can  be  held  in  regard  to  these  oxides  is  that 
in  them  the  nitrogen  is  univalent,  bivalent,  trivalent, 
quadrivalent,  and  quinquivalent,  a  view  which  is  ex- 
pressed by  the  following  formulas  : 

^  N=O  0=N=0 

£>0,    N=0,          0,     0=N=0, 


These    formulas  are,  however,  purely  speculative  and 
represent  nothing  known  to  us.      But  if  the  valence  of 


290  INORGANIC  CHEMISTRY. 

nitrogen  can  vary  in  this  way,  we  may  also  conceive  that 
the  oxygen  is  univalent  in  all  the  compounds  except 
nitrous  oxide.  Thus  nitric  oxide  may  be  represented  by 

the  formula  N-O,  nitrogen  peroxide  by  N  <  Q,  etc.     On 

the  other  hand,  there  is  an  unmistakable  tendency  on 
the  part  of  the  elements  to  act  either  with  even  valences, 
as  2,  4,  6,  etc.,  or  with  odd,  as  1,  3,  5,  etc.  This  is  beauti- 
fully shown  by  the  members  of  the  chlorine  group  and 
those  of  the  sulphur  group.  It  has  been  pointed  out 
that  the  relations  between  the  compounds  of  chlorine, 
bromine,  and  iodine  can  be  explained,  by  assuming  that 
these  elements  are  univalent,  trivalent,  quinquivalent, 
and  septivalent ;  and  that  the  relations  between  the 
compounds  of  sulphur,  selenium,  and  tellurium  can  be 
equally  easily  explained  by  assuming  that  these  elements 
are  bivalent,  quadrivalent,  and  sexivalent.  In  the  case 
of  nitrogen  and  the  elements  belonging  to  the  same  group 
we  should  naturally  expect  to  find  a  similar  law  of  com- 
position holding  good.  As  far  as  the  hydroxyl  deriva- 
tives, represented  by  nitrous  acid  and  nitric  acid,  are 
concerned,  the  same  regularity  is  observed  as  in  the  case 
of  sulphur.  In  nitric  acid  the  nitrogen  is  probably  quin- 
quivalent, and  in  nitrous  acid  trivalent.  Further,  in 
ammonia  nitrogen  is  trivalent,  while  it  is  probably  quin- 
quivalent in  the  ammonium  compounds,  as  has  been 
pointed  out  (see  p.  275).  It  is  clear  that  nitrogen  tends 
to  act  either  as  a  trivalent  or  quinquivalent  element. 
"Whether  it  ever  acts  as  a  univalent  element  it  is  impos- 
sible to  say,  for,  while  the  existence  of  the  compound 
N2O  seems  to  show  that  it  does,  this  same  compound 
may  be  explained  on  the  assumption  that  111  it  the  ni- 

N 
trogen  is  trivalent,  as  shown  in  the  formula  ||  >O;  and 

N 

indeed  there  is  no  difficulty  in  assuming  any  desired 
valence  for  the  nitrogen.  Taking  the  compound  nitric 
oxide,  there  seems  to  be  no  escape  here  from  the  con- 
clusion that  the  nitrogen  is  bivalent  if  the  oxygen  is  bi- 
valent ;  and  the  compound  forms  a  striking  exception  to 


STRUCTURE  OF  COMPOUNDS  OF  NITROGEN.        291 

the  rule  above  referred  to  that  the  valence  of  an  element 
generally  changes  from  odd  to  odd  or  from  even  to  even. 
It  may  be  said  that  this  compound  is  unsaturated,  and 
that  one  of  its  bonds  is  unemployed,  a  condition  which 
may  be  symbolized  by  this  expression,  -N= O,  but  this 
does  not  help  us  out  of  the  difficulty,  and,  further,  this 
conception  is  not  in  accordance  with  the  fact  that  nitric 
oxide  takes  up  one  atom  of  oxygen  to  form  nitrogen  per- 
oxide, NO2.  And  then  the  question  arises,  What  is  the 
structure  of  this  last-mentioned  compound  ?  Should  it 
be  represented  thus  :  O=N=O?  If  so  the  nitrogen  is 
quadrivalent.  But  it  passes  readily  into  the  form  N2O4. 
It  may  be  that  this  act  consists  simply  in  the  union  of 
the  two  molecules  by  means  of  the  fifth  bond  of  quin- 
quivalent nitrogen,  the  structure  of  the  resulting  mole- 
cule being  represented  thus :  I  .  All  this  is,  how- 
ever, almost  pure  speculation,  and,  at  the  present  stage 
of  our  knowledge  of  the  subject  of  structure,  the  above 
formulas  have  very  little  value.  Still  it  must  not  be 
forgotten  that  the  structure  of  all  chemical  compounds 
is  a  legitimate  subject  of  investigation. 

When  we  come  to  the  acids  of  nitrogen  it  is  seen,  as 
has  already  been  pointed  out,  that  these  can  be  explained 
very  satisfactorily  by  the  aid  of  the  same  hypothesis 
that  served  so  well  in  dealing  with  the  acids  of  iodine 
and  of  sulphur.  Nitric  acid  is  to  be  regarded  as  derived 
from  the  maximum  hydroxyl  compound  of  quinquivalent 
nitrogen,  known  as  normal  nitric  acid,  by  loss  of  water ; 
and  in  a  similar  way  nitrous  acid  is  to  be  regarded  as 
derived  from  the  maximum  hydroxyl  compound  of  tri- 
valent  nitrogen,  or  normal  nitrous  acid,  by  loss  of  water. 
A  few  salts  are  known  which  appear  to  be  derived  from 
the  normal  acids,  but  for  the  most  part  all  the  hydrogen 
atoms  of  these  normal  acids  are  not  replaceable  by 
metals,  and  the  formation  of  salts  generally  involves  a 
breaking  down  of  the  compound  into  water  and  the  com- 
mon form  of  the  acid. 

Compounds  of  Nitrogen  with  the  Elements  of  the  Chlo- 
rine Group. — Notwithstanding  the  ease  with  which  chlo- 


292  INORGANIC  CHEMISTRY. 

rine  combines  with  most  elements,  and  the  stability  of 
the  compounds  which  it  forms  with  them,  its  compound 
with  nitrogen  is  extremely  unstable.  It  can  be  made  by 
the  action  of  chlorine  on  ammonia,  and  by  decomposing 
a  solution  of  ammonium  chloride  by  means  of  an  electric 
current.  In  the  latter  case  chlorine  is  liberated  at  one 
of  the  poles  and  then  acts  upon  the  ammonium  chloride  : 

NH4C1  +  601  =  4HC1  +  NCI,. 

It  appears  that  when  chlorine  acts  upon  ammonia  differ- 
ent products  are  formed  by  successive  replacement  of 
the  hydrogen  atoms  of  the  ammonia  by  chlorine,  thus : 

NH3      +  C12  =  NH,C1  +  HC1 ; 
NH2C1  +  Cl,  =  NHC12  +  HC1 ; 

NHCla  +  C12  =  NC13     +  HC1. 

According  to  this,  the  trichloride  of  nitrogen  is  the  final 
product  of  the  substituting  action  of  chlorine  upon  am- 
monia. The  compound  is  an  oil,  which  undergoes  de- 
composition very  readily.  It  is,  indeed,  one  of  the  most 
explosive  substances  known.  It  is  decomposed  by  heat, 
and  especially  by  contact  with  certain  substances,  among 
which  are  oil  of  turpentine  and  caoutchouc.  It  is  slowly 
decomposed  by  water,  though,  probably  owing  to  the 
slight  affinity  of  nitrogen  for  oxygen,  the  decomposition 
does  not  take  place  as  readily  as  that  of  the  compounds 
of  sulphur  and  chlorine.  Direct  sunlight  causes  explo- 
sion of  the  chloride. 

When  ammonia  is  treated  with  iodine  reactions  take 
place  similar  to  those  which  take  place  with  chlorine. 
The  products  are  the  iodides  of  nitrogen,  the  final 
product  of  the  action  being  the  tri-iodide,  NI3.  These 
compounds,  like  the  corresponding  chlorine  compounds, 
are  extremely  explosive.  The  simplest  way  to  prepare 
them  is  to  place  a  little  powdered  iodine  on  a  filter  and 
pour  concentrated  ammonia  over  it.  The  substance 
should  be  made  in  only  very  small  quantities  at  a  time. 
"When  dried  it  decomposes  with  violent  explosion  by 
contact  even  with  soft  substances ;  and  it  will  also  ex- 


COMPOUNDS  OF  NITROGEN  WITH  SULPHUR,  ETC.    293 

plode  if  left  entirely  undisturbed.  The  different  com- 
pounds called  nitrogen  iodide  are  slowly  decomposed  by 
water. 

Compounds  of  Nitrogen  with  the  Members  of  the  Sul- 
phur Group.  —  Nitrogen  combines  with  sulphur  forming 
two  compounds,  N4S4  and  N6S2,  both  of  which  are  well 
characterized.  Among  the  most  interesting  compounds 
containing  sulphur  and  nitrogen  is  that  which  has  been 
referred  to  as  nitrosyl-sulphuric  acid  in  connection  with 
the  account  of  the  manufacture  of  sulphuric  acid.  It  is 
formed  by  the  action  of  sulphur  dioxide  on  fuming  nitric 
acid  : 


also  in  the  manufacture  of  sulphuric  acid  by  the  action 
of  sulphur  dioxide,  water,  and  oxygen  upon  a  mixture 
of  nitrogen  peroxide  and  nitric  oxide  : 

280,  +  H.O  +  N,0,  +  20  =  280,<gg>; 

and  by  the  action  of  the  nitrogen  peroxide  upon  sul- 
phuric acid  : 


2N02  +  S0g<        =  80,<»  +HNOa. 


When  treated  with  water  it  breaks  down  into  sulphuric 
acid  and  nitrogen  trioxide.  It  will  be  remembered  that, 
in  order  to  prevent  loss  of  oxides  of  nitrogen  in  the  sul- 
phuric acid  factories,  the  gases  are  brought  in  contact 
with  concentrated  sulphuric  acid  in  the  Gay  Lussac  tower 
before  being  allowed  to  escape.  The  oxides  form  with 
sulphuric  acid  compounds  similar  to  nitrosyl-sulphuric 
acid,  and  when  these  are  diluted  with  "  chamber  acid  " 
and  heated  by  the  hot  gases  from  the  sulphur  furnace 
the  oxides  of  nitrogen  are  given  up. 


CHAPTER  XVII. 

ELEMENTS   OF   FAMILY  V,  GROUP   B : 

PHOSPHORUS— ARSENIC— ANTIMONY— BISMUTH. 

THE   ELEMENTS    AND    THEIR    COMPOUNDS    WITH 

HYDROGEN. 

General. — The  elements  of  this  group  bear  to  nitrogen 
very  much  the  same  relations  that  the  members  of  the 
sulphur  group  bear  to  oxygen,  and  those  of  the  chlorine 
group  bear  to  fluorine.  In  general  they  form  compounds 
of  the  same  character  and  of  similar  composition.  At 
the  same  time  gradations  in  properties  are  noticed  in 
passing  from  one  end  of  the  group  to  the  other.  Like 
nitrogen,  the  elements  of  the  group  are  strongly  marked 
acid-formers,  though  this  character  grows  less  marked 
from  nitrogen  to  bismuth.  Antimony  is  both  an  acid- 
forming  and  a  base-forming  element,  while  bismuth  is 
more  basic  than  acid.  The  stability  of  the  hydrogen 
compounds  decreases  from  nitrogen  to  antimony  ;  while 
bismuth  does  not  form  a  compound  with  hydrogen. 
Ammonia,  as  we  have  s"een,  is  strongly  basic ;  the  cor- 
responding compound  of  phosphorus  and  hydrogen  has 
weak  basic  properties,  while  those  of  arsenic  and  anti- 
mony have  no  basic  properties.  These  hydrogen  com- 
pounds correspond  in  composition  to  ammonia.  They 
are : 

NH3:        PH3        AsH3        SbH3 

With  chlorine  they  all  form  compounds  corresponding 
to  nitrogen  trichloride,  and  phosphorus  and  antimony 
form  compounds  in  which  they  are  quinquivalent,  while 

(294) 


ELEMENTS  OF  FAMILY  V,    GROUP  B.  295 

bismuth  forms  a  chloride,  Bi2Cl4 ,  in  addition  to  the  tri- 
chloride. The  compounds  referred  to  are  : 

Bi2Cl4 

NC13:        PC13        AsCl3        SbCls        BiCl3 
PC15  SbCl6 

They  all  form  two  oxides  corresponding  to  nitrogen  tri- 
oxide  and  pentoxide : 

N203:        P203        As203        Sb203        Bi2O3 
N206:        P205        As  A        Sb205        BiA 

No  one  of  the  elements  of  the  group  forms  as  great  a 
variety  of  compounds  with  oxygen  as  nitrogen  does. 
Antimony,  however,  forms  the  oxide  Sb2O4,  correspond- 
ing to  nitrogen  peroxide,  N  2O4 ;  and  bismuth  forms  the 
oxide  Bi2O2  or  BiO,  corresponding  to  nitric  oxide,  NO. 

The  hydroxyl  compounds  or  acids,  like  those  of  nitro- 
gen, are  related  to  the  maximum  hydroxyl  compounds 
of  the  elements  with  the  valence  5,  and  to  the  maximum 
hydroxyl  compounds  of  the  elements  with  the  valence  3. 
That  is  to  say,  they  may  be  regarded  as  derived  from  a 
hydroxide  of  the  general  formula  M(OH)B,  and  another 
of  the  formula  M(OH)3.  Where  M  is  nitrogen  these  acids 
break  down  to  the  forms  NO2(OH)  and  NO(OH)  by  loss 
of  one  or  two  molecules  of  water.  In  the  case  of  the 
elements  of  the  phosphorus  group,  however,  the  breaking 
down  is  not  generally  carried  as  far  as  with  nitrogen. 
The  general  rule  is  the  same  as  in  the  sulphur  and  chlo- 
rine groups  :  the  normal  acid  breaks  down  to  form  com- 
pounds containing  the  same  number  of  hydrogen  atoms 
as  the  hydrogen  compounds  of  the  elements.  Thus  the 
hydroxyl  derivatives  of  chlorine  generally  break  down  to 
form  compounds  containing  one  atom  of  hydrogen,  or  the 
same  number  that  is  contained  in  the  hydrogen  compound, 
hydrochloric  acid,  thus:  Cl(OH),  yields  C1O3(OH) ; 
C1(OH)6  yields  GLOa(OH),  etc.  So,  also,  in  the  sulphur 
group,  S(OH)6  yields  SO9(OH),,  etc.,  the  number  of  hy- 
drogen atoms  in  the  common  form  of  the  acid  being  the 
same  as  that  in  the  hydrogen  compound  of  sulphur,  SHa. 


296  INORGANIC  CHEMISTRY. 

This  rule  does  not  hold  good  for  nitrogen,  for  the 
tendency  here  is  to  break  down  to  compounds  contain- 
ing one  atom  of  hydrogen.  On  the  other  hand,  phos- 
phorus, arsenic,  and  antimony  follow  the  rule,  the 
principal  acids  of  these  elements  containing  three  atoms 
of  hydrogen  in  the  molecule.  As  already  stated,  there 
are  two  series  of  these  represented  by  the  following  for- 
mulas : 

H3P03        H3As03        H3Sb03 
H3PO4        H3AsO4        H3SbO4 

Besides  the  above,  however,  phosphorus  forms  several 
other  acids.  The  principal  ones  bear  simple  relations  to 
the  acid  H3PO4,  which  is  called  orthophosphoric  acid. 
The  simplest  view  in  regard  to  the  acid  of  phosphorus 
of  the  formula  H3PO4,  and  the  corresponding  compounds 
of  arsenic  and  antimony,  is  that  it  is  derived  from  the 
corresponding  normal  acid  by  loss  of  one  molecule  of 
water.  Thus,  normal  phosphoric  acid  is  P(OH)5.  By 
loss  of  one  molecule  of  water  this  yields  ordinary  or 
orthophosphoric  acid : 

P(OH)5  =  PO(OH)3  +  H20. 

Normal  phosphorous  acid  is  P(OH)3.  Whether  or- 
dinary phosphorous  acid  has  this  structure  is  a  ques- 
tion very  difficult  to  answer  at  present.  By  loss  of 
another  molecule  of  water  orthophosphoric  acid  is  con- 
verted into  metaphosphoric  acid,  which  in  composition 
corresponds  to  nitric  acid.  Its  formation  is  represented 
thus : 

PO(OH)3  =  P02(OH)  +  H20. 

The  series  of  phosphorus  acids,  H3PO2,  H3PO3,  and 
H3PO4  is  strongly  suggestive  of  the  series  of  sulphur 
acids,  H2SO2,  H2SO3,  and  H2SO4,  and  of  the  series  of 
chlorine  acids,  HC1O,  HC1O2,  HC1O3,  and  HC1O4. 

Arsenic  and  antimony  also  form  acids  corresponding 
to  metaphosphoric  acid,  known  respectively  as  metarsenic 
and  metantimonic  acids. 


ELEMENTS  OF  FAMILY  V,    GROUP  B.  297 

By  elimination  of  one  molecule  of  water  from  two 
molecules  of  ordinary  phosphoric  acid  there  is  formed 
an  acid  H4P2O,,  known  as  pyrophosphoric  acid,  which 
bears  to  ordinary  phosphoric  acid  much  the  same  rela- 
tion that  pyrosulphuric  or  disulphuric  acid  bears  to 
ordinary  sulphuric  acid  : 

/OH 
SO/ 

>0     +  H30. 


/OH 
PO-OH 


=      \0         H  0 

/OH         PO/OH 
^-(J±l  \rm" 

\OH 

Ordinary  arsenic  and  antimonic  acids  yield  corre- 
sponding derivatives  known  as  pyroarsenic  and  pyroanti- 
monic  acids.  The  elements  of  the  phosphorus  group 
form  compounds  with  oxygen  and  chlorine  known  as  the 
oxy  chlorides,  which  in  general  resemble  the  oxychlorides 
of  the  members  of  the  sulphur  group.  Examples  of  these 
compounds  are  phosphorus  oxychloride,  POC13,  anti- 
mony oxychloride,  SbOCl,  and  bismuth  oxychloride, 
BiOCl.  Phosphorus  oxychloride  is  readily  decom- 
posed by  water,  forming  phosphoric  and  hydrochlo- 
ric acids  : 

(  Cl        HOH  (  OH 

PO^  Cl  +  HOH  =  PO^  OH  +  3HC1. 
(  Cl        HOH  (  OH 

The  oxychlorides  of  antimony  and  bismuth  are  not 
completely  decomposed  by  water.  This  is  in  accordance 
with  the  fact  to  which  attention  has  been  called  that  the 
chlorides  of  the  acid-forming  elements  are  in  general 
easily  decomposed  by  water  and  converted  into  hydroxyl 
compounds,  while  the  chlorides  of  the  base-forming  ele- 
ments are  not  readily  decomposed  in  this  way,  but,  on 
the  contrary,  their  oxides  and  hydroxides  are,  as  a  rule, 


298  INORGANIC  CHEMISTRY. 

readily  converted  into  chlorides  by  hydrocliloric  acid 
Elements  which,  like  antimony  and  bismuth,  play  the 
part  of  base-formers  and  acid-formers  form  stable 
oxy  chlorides. 

Of  the  elements  of  this  group  phosphorus  occurs 
most  abundantly  in  nature,  arsenic  and  antimony  next, 
and  bismuth  least  abundantly.  Arsenic  and  bismuth 
occur  to  some  extent  in  the  uncombined  condition. 
Phosphorus  and  antimony  occur  in  combination.  All 
the  elements  of  the  group  find  applications  in  the  arts, 
either  as  the  elements  or  in  the  form  of  compounds. 

PHOSPHORUS,  P  (At.  Wt.  30.79). 

• 

Occurrence. — The  name  phosphorus  is  derived  from 
the  Greek  0c3?,  light,  and  (pspeiv,  to  carry,  on  account 
of  the  fact  that  it  always  gives  light  and  takes  fire  very 
easily.  The  element  occurs  in  nature  in  the  form  of 
phosphates  derived  from  orthophosphoric  acid,  H3PO4. 
The  chief  of  these  is  calcium  phosphate,  Ca3(PO4)2,  which 
is  the  principal  constituent  of  the  minerals  phosphorite 
and  apatite,  and  of  the  ashes  of  bones.  The  phosphates, 
like  the  nitrates,  are  widely  distributed  in  the  soil  and 
are  of  fundamental  importance  in  the  process  of  plant 
life.  The  phosphates  found  in  the  bones  are  taken  into 
the  animal  body  in  the  food.  All  plants  used  as  food 
contain  small  quantities  of  the  phosphates  which  they 
get  from  the  soil.  The  phosphates  taken  into  the  body 
are  partly  given  off  in  the  excrement  and  urine,  and  it 
was  in  an  examination  of  urine  made  in  the  hope  of 
finding  the  philosopher's  stone  that  phosphorus  was 
first  discovered  in  1669.  At  present  phosphorus  is 
made  almost  entirely  from  bones. 

Preparation. — Besides  the  phosphates,  considerable 
quantities  of  organic  materials  are  contained  in  bones. 
When  the  bones  are  burned  the  organic  materials  pass 
off  for  the  most  part  in  the  form  of  carbon  dioxide,  water, 
and  volatile  compounds  containing  nitrogen,  and  the  so- 
called  mineral  or  earthy  portions,  the  chief  constituent 
of  which  is  tertiary  calcium  phosphate,  or  phosphoric 


PHOSPHORUS:— OCCURRENCE- PREPARATION.      299 

acid  in  which  all  the  hydrogen  is  replaced  by  calcium, 
remain  behind.  As  calcium  is  bivalent  and  there  are 
three  atoms  of  hydrogen  in  the  molecule  of  phosphoric 
acid,  H3PO4,  the  simplest  way  in  which  all  the  hydrogen 
atoms  of  the  acid  can  be  replaced  by  calcium  is  that 
represented  by  the  formula  Ca3(PO4)2,  the  six  atoms  of 
hydrogen  in  two  molecules  of  the  acid  being  replaced  by 
three  bivalent  atoms  of  calcium.  The  problem  now  is 
to  isolate  the  phosphorus  from  this  calcium  phosphate. 
The  salt  is  insoluble  in  water,  and  there  is  no  simple 
way  by  which  the  phosphorus  can  be  set  free  from  it. 
When  it  is  treated  with  sulphuric  acid  calcium  sulphate 
which  is  difficultly  soluble  is  deposited  and  phosphoric 
acid  is  set  free  and  remains  in  solution.  The  reaction  is 
represented  as  follows  : 

Ca3(P04)2  +  3H2SO4  =  2H,POt  +  3CaSO4. 

The  calcium  sulphate,  or  gypsum,  is  allowed  to  settle 
and  is  then  filtered  off  and  washed.  The  solution  of 
phosphoric  acid  is  evaporated  until  it  has  the  specific 
gravity  1.325  to  1.5.  Then  it  is  mixed  with  coarsely- 
ground  wood  charcoal,  coke,  or  sawdust,  and  carefully 
dried  in  a  cast-iron  pot  or  muffle  furnace.  In  this  pro- 
cess the  orthophosphoric  acid,  H8PO4 ,  is  converted  into 
metaphosphoric  acid,  HPO3, 

H3P04  =  HP03  +  H20 ; 

and  in  case  sawdust  is  used  this  is  changed,  so  that 
the  resulting  mixture  consists  of  charcoal  and  metaphos- 
phoric acid.  This  mixture  is  next  subjected  to  distilla- 
tion in  clay  retorts,  when  the  metaphosphoric  acid  is  re- 
duced according  to  the  following  equation : 

2HP03  +  60  =  H2  +  6CO  +  2P. 

The  phosphorus  passes  over  in  the  form  of  vapor,  and 
is  collected  under  water.  The  crude  phosphorus  thus 
obtained  must  be  subjected  to  a  cleansing  process  before 
it  can  be  used.  For  this  purpose  it  is  pressed,  while  in 
the  molten  condition  under  water,  through  chamois 
leather,  or  it  is  distilled  again  from  iron  retorts  ;  or,  still 
better,  it  is  treated  with  chromic  acid  as  follows :  It  is 
fused  under  water,  then  a  little  potassium  or  sodium 


300  INORGANIC  CHEMISTRY. 

bichromate  in  solution  is  added,  and  afterwards  an  equiv- 
alent proportion  of  sulphuric  acid,  and  the  whole  al- 
lowed to  stand  for  two  hours  or  more.  The  phosphorus 
is  then  washed  with  hot  water,  and  after  being  si- 
phoned off  it  is  filtered  through  canvas  bags.  The 
phosphorus  is  then  cast  into  sticks  in  tin  tubes.  In  this 
form  it  generally  comes  into  the  market. 

At  the  time  of  the  last  report  available  there  were 
manufactured  in  one  year  about  1200  tons  of  phospho- 
rus in  two  factories,  one  of  which  is  in  England  and  the 
other  in  France.  Quite  recently  phosphorus  has  been 
manufactured  to  some  extent  in  Sweden. 

Properties. — Ordinary  phosphorus  is  colorless  or 
slightly  yellowish,  translucent,  and  at  ordinary  tempera- 
tures it  can  be  cut  like  wax,  but  it  becomes  hard  and 
brittle  at  low  temperatures.  It  melts  at  44°,  and  boils 
at  290°.  It  is  insoluble  in  water.  When  kept  under 
water  for  any  length  of  time  in  dispersed  light  it  be- 
comes opaque,  crystalline  on  the  surface,  and  yellow. 
It  is  soluble  in  carbon  disulphide,  and  crystallizes  when 
deposited  from  this  solution.  It  gives  off  fumes  in  con- 
tact with  the  air,  and  emits  a  pale  light  which  is  known 
as  a  phosphorescent  light.  It  is  very  poisonous,  the  in- 
halation of  the  vapor  in  small  quantities  causing  very 
serious  disturbance  of  the  system.  The  workmen  in  the 
factories  where  phosphorus  is  made  or  used  are  fre- 
quently affected  by  phosphorus-poisoning.  Among  the 
prominent  symptoms  is  gradual  decomposition  of  the 
bones.  When  taken  into  the  stomach  phosphorus  also 
acts  as  a  poison  and  causes  death.  When  heated  in  the 
air  it  takes  fire  at  50°.  It  also  takes  fire  by  rubbing, 
and  it  must  be  handled  with  the  greatest  care,  as  wounds 
caused  by  it  are  dangerous  and  difficult  to  heal.  When 
it  burns  in  the  air  it  is  converted  into  the  pentoxide, 
P2O5,  which  is  also  the  product  of  its  combustion  in 
oxygen,  as  we  have  seen.  It  combines  also  with  other 
elements  directly,  frequently  with  evolution  of  light. 
Thus,  when  it  is  brought  together  with  chlorine,  bromine, 
and  iodine,  it  forms  the  compounds  PC13,  PBr3,  and  PI3. 
It  also  combines  with  sulphur.  When  a  piece  is  put  in 
water  and  the  water  boiled,  a  part  of  the  phosphorus 


PROPERTIES  OF  PHOSPHORUS.  301 

passes  over,  and  if  the  water  vapor  is  condensed  in  a 
glass  tube  in  a  dark  room,  it  is  seen  to  be  phosphores- 
cent. This  furnishes  a  convenient  method  for  its  detec- 
tion, as,  for  example,  in  a  case  of  suspected  poisoning  by 
phosphorus. 

Owing  to  its  strong  tendency  to  combine  with  oxygen, 
it  abstracts  the  element  from  some  of  its  compounds. 
Thus,  if  a  solution  of  phosphorus  in  carbon  disulphide 
is  added  to  a  solution  of  copper  sulphate,  metallic  cop- 
per is  thrown  down,  while  at  the  same  time  copper 
phosphate  and  a  compound  of  copper  and  phosphorus 
are  formed. 

When  phosphorus  is  left  for  a  long  time  under  water 
and  subjected  to  the  action  of  light,  it  becomes  at  first 
yellow,  then  reddish,  and  finally  red.  The  same  change 
takes  place  when  phosphorus  is  heated  for  a  time  in  an 
atmosphere  which  is  free  from  oxygen;  and  rapidly 
when  it  is  heated  to  300°  in  an  hermetically  sealed  tube. 
The  red  substance  thus  obtained  has  properties  entirely 
different  from  those  of  ordinary  phosphorus.  It  is  a 
red  powder,  which  frequently  has  a  crystalline  struc- 
ture. It  does  not  emit  light.  It  does  not  melt  at  a  low 
temperature.  It  is  not  poisonous,  and  cannot  be  easily 
ignited.  Further,  it  is  perfectly  insoluble  in  carbon 
disulphide.  In  every  respect  this  red  modification  of 
phosphorus  conducts  itself  as  a  much  less  active  sub- 
stance chemically  than  ordinary  phosphorus.  In  an 
atmosphere  of  carbon  dioxide  it  is  converted  into  ordi- 
ary  phosphorus  when  heated  to  261°,  and  if  heated  to 
this  temperature  in  the  air  it  takes  fire,  and  then  forms 
the  same  product  that  ordinary  phosphorus  does  in 
burning. 

When  phosphorus  is  heated  with  lead  for  eight  to  ten 
hours  to  a  very  high  temperature  in  sealed  tubes  from 
which  the  air  has  been  exhausted,  and  the  whole  then 
allowed  to  cool,  the  surface  of  the  lead  is  found  covered 
with  black,  laminated  crystals,  which  undergo  no  change 
in  the  air.  Crystals  are  also  found  in  the  interior  of  the 
lead.  This  variety  of  phosphorus  is  called  crystallized, 
metallic  phosphorus  on  account  of  the  metallic  lustre.  It 
is  not  as  volatile  as  the  ordinary  variety. 


302  INORGANIC  CHEMISTRY. 

When  the  vapor  of  phosphorus  is  suddenly  cooled  by 
ice  water  in  an  atmosphere  of  hydrogen,  it  is  deposited 
in  the  form  of  a  snow-white  powder  on  the  surface  of 
the  water.  Under  water  this  variety  undergoes  very 
little  change  even  when  exposed  for  a  long  time  to  the 
action  of  the  sunlight.  When  exposed  to  the  air  on 
filter-paper  it  gives  off  dense  fumes,  and  then  melts, 
forming  ordinary  phosphorus,  but  it  does  not  generally 
take  fire  under  these  circumstances. 

Treated  with  oxidizing  agents,  as,  for  example,  nitric 
acid,  phosphorus  is  slowly  converted  into  phosphoric 
acid,  just  as  sulphur  is  converted  into  sulphuric  acid 
under  the  same  conditions. 

Applications  of  Phosphorus.-  —  Phosphorus  is  used  prin- 
cipally in  the  manufacture  of  matches  and  as  a  poison 
for  vermin.  Various  mixtures  are  used  for  making 
matches.  Nearly  all  of  them  contain  phosphorus  to- 
gether with  some  oxidizing  compound,  and  some  neutral 
substance  to  act  as  a  medium  for  holding  the  constitu- 
ents together.  An  example  is  a  mixture  consisting  of  2 
parts  phosphorus,  1  part  manganese  dioxide,  3  parts 
chalk,  -J  part  lamp-black,  and  5  parts  glue.  The  mix- 
ture used  in  the  manufacture  of  the  so-called  "  safety 
matches"  consists  of  potassium  chlorate,  potassium 
dichromate,  minium,  and  antimony  trisulphide.  This 
will  not  ignite  by  simple  friction,  but  will  ignite  when 
drawn  across  a  paper  upon  which  is  a  mixture  of  red 
phosphorus  and  antimony  pentasulphide. 

Compounds  of  Phosphorus  with  Hydrogen.  —  There  are 
three  compounds  of  phosphorus  with  hydrogen,  a  gaseous 
compound  of  the  formula  PH3,  corresponding  to  am- 
monia ;  a  liquid  of  the  formula  PH2  ,  or  PQH4  ,  correspond- 
ing to  hydrazine  ;  and  a  solid  of  the  formula  P,H3,  or  P4Ha. 

Phosphine,  Gaseous  Phosphuretted  Hydrogen,  PH3.  — 
This  compound  is  formed  when  phosphorous  or  hypo- 
phosphorous  acid  is  heated.  The  decompositions  take 
place  as  represented  in  these  equations  : 


4H3P03:=3H3P04 

2=    H3P04-fPH8. 


PHOSPHINE.  303 

"We  see  here  an  example  of  the  same  kind  of  action  that 
was  referred  to  in  connection  with  the  sulphur  com- 
pounds. It  will  be  remembered  that,  in  general,  when 
a  salt  of  any  oxygen  acid  of  sulphur  except  sulphuric 
acid  is  heated  it  is  converted  into  the  sulphate,  and  that 
the  other  elements  arrange  themselves  in  simpler  forms 
of  combination.  Thus,  when  sodium  thiosulphate  is 
heated  it  is  converted  into  sodium  sulphate  and  sodium 
pentasulphide,  as  represented  in  the  following  equation : 

4Na2S2O3  =  3Na2SO4  +  Na2S5. 

So  also  sodium  sulphite  yields  sodium  sulphate  and 
sodium  sulphide : 

4Na2SO8  =  3Na2SO4  +  NaaS. 

Other  ways  of  making  phosphine  are  :  (1)  By  treating 
a  strong  solution  of  potassium  hydroxide  with  phos- 
phorus, when  reaction  takes  place  as  follows : 

3KOH  +  4P  +  3H20  =  3KH2PO2  +  PH3. 

The  compound  KH2PO2  is  known  as  potassium  hypo- 
phosphite,  being  derived  from  hypophosphorous  acid, 
H3PO.2.  (2)  By  treating  zinc  phosphide  with  dilute 
hydrochloric  acid.  Assuming  that  zinc  phosphide  has 
the  composition  represented  by  the  formula  Zn3P2,  the 
reaction  with  hydrochloric  acid  takes  place  according  to 
the  equation 

Zn3P2  +  6HC1  =  3ZnCl2  +  2PH3. 

(3)  By  treating  phosphonium  iodide,  PH4I,  with  water  or 
a  dilute  solution  of  potassium  hydroxide : 

PH4I  +  H20    =PH3  +  HI  +  H2O; 

PHJ  +  KOH  =  PH3  +  KI  +  H20. 

When  made  from  phosphorus  and  potassium  hydroxide 
it  always  contains  a  considerable  proportion  of  hydrogen, 
for  the  reason  that  potassium  hypophosphite  gives  off 
hydrogen  when  heated  with  a  solution  of  potassium 
hydroxide.  From  calcium  phosphide  and  from  phos- 
phonium iodide  it  can  be  obtained  in  pure  condition. 


304  INORGANIC  CHEMISTRY. 

Phosphine  is  a  colorless  gas  with  an  unpleasant,  gar- 
lic-like odor.  It  is  insoluble  in  water,  and  is  poisonous. 
It  burns,  but  does  not  take  fire  spontaneously  when 
pure.  When  burned  with  free  access  of  air  the  products 
of  combustion  are  phosphorus  pentoxide  and  water  : 


whereas  when  it  is  burned  in  a  cylinder  so  that  the  air 
has  not  free  access  to  it,  the  products  are  water  and 
phosphorus,  which  is  deposited  in  a  reddish  layer  upon 
the  glass. 

Although  pure  phosphine  does  not  take  fire  spontane- 
ously when  brought  in  contact  with  the  air,  the  gas 
made  by  any  one  of  the  methods  above  referred  to  is 
pretty  sure  to  contain  some  of  the  liquid  compound  of 
phosphorus  and  hydrogen,  P2H4,  which  is  spontaneously 
inflammable,  and  therefore  the  gas  takes  fire.  If  it  is 
collected  in  a  glass  vessel  over  water,  and  allowed  to 
stand  so  that  the  light  acts  upon  it,  the  liquid  phosphine 
is  decomposed  into  the  gaseous  and  solid  varieties,  and 
the  gas  which  is  left  no  longer  has  the  property  of  tak- 
ing fire  spontaneously.  Phosphine  is  much  less  stable 
than  ammonia.  When  heated  or  when  treated  with  elec- 
tric sparks  it  is  easily  decomposed  into  phosphorus  and 
hydrogen.  While  ammonia  dissolves  in  water,  probably 
forming  the  hydroxide  NH4(OH),  phosphine  is  only  very 
slightly  soluble  in  water.  Ammonia  combines  with  acids 
very  energetically,  forming  the  ammonium  salts,  and  we 
should  expect  to  find  that  similar  salts  are  formed  by 
the  action  of  phosphine  on  acids  ;  but  only  a  few  such 
compounds  are  known,  and  these  are  unstable.  Thus, 
when  phosphine  is  brought  together  with  hydrochloric, 
hydrobromic,  and  hydriodic  acids,  the  reactions  repre- 
sented by  the  following  equations  take  place  : 

PH3  +  HC1  =  PH4C1  ; 
PH3  +  HBr  =  PH4Br  ; 
PH3  +  HI  =PH4I. 

The  products  are  called  respectively  phosphonium  Mo- 
ride,  bromide,  and  iodide.  The  reactions  are,  as  will  be 


ARSENIC :— OCCURRENCE—  PREPARATION.  305 

seen,  perfectly  analogous  to  those  which  take  place  be- 
tween the  same  acids  and  ammonia.  But  the  products 
are  much  less  stable  than  the  ammonium  salts.  The 
bromide  when  exposed  to  the  air  attracts  water  and 
decomposes  rapidly,  forming  hydrobromic  acid  and 
phosphine.  Phosphonium  iodide  undergoes  a  similar 
decomposition. 

AESENIC,  As  (At.  Wt.  74.44). 

Occurrence. — Arsenic  occurs  in  nature  to  some  extent 
in  the  uncombined  condition  or  native.  Compounds  of 
the  metals  with  arsenic,  or  the  arsenides,  occur  very 
widely  distributed,  and  they  frequently  accompany,  and 
are  similar  to,  the  sulphides.  The  most  common  com- 
pound of  this  kind  is  the  so-called  arsenical  pyrites, 
which  has  the  composition  FeAsS,  and  may  therefore 
be  regarded  as  iron  pyrites,  FeS2,  in  which  one  atom  of 
sulphur  has  been  replaced  by  one  atom  of  arsenic. 
Among  other  arsenic  compounds  deserving  special  men- 
tion are  the  two  arsenides  of  iron  of  the  formulas  FeAs2 
and  Fe2As3,  which  are  apparently  analogous  to  the  sul- 
phides FeS2  and  Fe2S3;  and,  further,  the  sulphides  of 
arsenic,  orpiment,  As2S3,  and  realgar,  As2S2.  The  oxide 
As2O3  occurs  in  considerable  quantity,  and  also  salts  of 
arsenic  acid,  or  the  arsenates,  which  in  composition  are 
analogous  to  the  phosphates. 

Preparation. — The  arsenic  which  conies  into  the  market 
is  either  that  which  occurs  native  or  it  is  made  from 
arsenical  pyrites  by  heating  : 

FeAsS  =  FeS  +  As. 

The  arsenic  thus  obtained  is  not  pure.  By  bringing  a 
little  iodine  in  the  bottom  of  a  porcelain  crucible,  put- 
ting the  arsenic  upon  it,  and  heating,  the  arsenic  ac- 
quires a  high  metallic  lustre,  and  once  in  this  condition 
it  will  remain  so  for  some  time  even  when  exposed  to 
the  air. 

Properties. — Arsenic  has  a  metallic  lustre  and  steel 
color.  It  is  very  brittle.  When  heated  it  volatilizes 


306  INORGANIC  CHEMISTRY. 

without  melting.  At  red  heat  it  burns  with  a  bluish 
flame,  and  the  vapor  given  off  has  the  odor  of  garlic. 
This  odor  produced  under  such  circumstances  is  very 
characteristic  of  arsenic,  and  furnishes  one  of  the  means 
for  detecting  it.  Arsenic  combines  with  most  elements 
directly,  the  action  being  accompanied  in  some  cases,  as 
in  that  of  chlorine,  by  an  evolution  of  light.  As  an  ele- 
ment it  is  not  poisonous,  but  when  oxidized  to  the  form 
of  the  oxide  As2O3  it  is  extremely  poisonous.  As  it  is 
easily  oxidized,  the  element  itself  may  act  as  a  poison. 
When  boiled  with  nitric  acid  arsenic  is  converted  into 
arsenic  acid,  H3AsO4,  just  as  phosphorus  is  converted  by 
nitric  acid  into  phosphoric  acid,  H3PO4. 

One  peculiarity  in  the  conduct  of  arsenic  is  suggestive, 
and  that  is  its  power  to  form  compounds  which  are  an- 
alogous to  the  compounds  of  sulphur.  There  are  a 
number  of  compounds  similar  to  arsenical  pyrites  which 
appear  to  be  perfectly  analogous  to  the  sulphur  com- 
pounds, and  in  them  it  seems  necessary  to  assume  that 
the  arsenic  plays  the  same  part  as  the  sulphur.  On  the 
other  hand,  arsenic  conducts  itself  in  nearly  all  its  com- 
pounds like  phosphorus.  This  power  to  play  double 
parts  is  not  uncommon  among  the  elements,  and  we  shall 
hereafter  meet  with  a  number  of  examples.  The  case  of 
manganese  is  one  in  point.  While  it  conducts  itself  in 
some  of  its  compounds  like  the  members  of  the  chlo- 
rine group,  to  which  on  account  of  its  position  in  the 
periodic  system  we  should  expect  to  find  it  related,  yet 
it  is  perhaps  more  closely  related  to  iron  and  chromium, 
which  belong  to  different  groups  ;  and  so,  also,  chromium, 
which  in  many  respects  resembles  sulphur  very  striking- 
ly, is  like  iron  and  aluminium  in  other  respects. 

Arsine,  Arseniuretted  Hydrogen,  AsH3. — This  com- 
pound is  analogous  to  ammonia  and  to  gaseous  phos- 
phine.  It  is  made  by  reduction  of  compounds  of  arsenic 
containing  oxygen,  as  arsenic  trioxide  or  arsenic  acid ; 
and  also  by  treating  a  compound  of  zinc  and  arsenic  with 
dilute  sulphuric  acid.  The  reactions  involved  in  the 
first  method  are 


ARSINE.  307 

As2O3     +  6H2  =  2AsH3  +  3H2O  ; 
H3AsO4  +  4H2  =  AsH3    +  4H2O. 

-  — --—_ 

That  involved  in  the  second  method  mentioned  is : 
As3Zn3  +  3H2SO4  =  2AsH3  +  3ZnSO4. 

Arsine  is  a  colorless  gas  with  an  odor  suggestive 
of  garlic.  It  is  extremely  poisonous,  even  very  small 
quantities  being  capable  of  producing  bad  effects,  and  it 
requires  but  little  to  cause  death.  When  ignited  in  the 
air  it  takes  fire  and  burns  with  a  pale  blue  flame,  the 
products  of  the  combustion  being  arsenic  trioxide,  As2O3, 
and  water.  If  the  air  is  prevented  from  gaining  free 
access  to  it  the  hydrogen  burns,  but  the  arsenic  is 
deposited  as  a  brownish  layer.  The  gas  is  so  unstable 
that,  when  it  is  passed  through  a  glass  tube  heated  to 
redness,  it  is  decomposed  into  arsenic  and  hydrogen,  the 
former  being  deposited  just  in  front  of  the  heated  por- 
tion of  the  tube  as  a  thin,  almost  black,  layer  with  a  high 
metallic  lustre. 

Arsine  is  easily  decomposed  by  most  active  chemical 
substances.  Water  and  concentrated  acids  decompose 
it ;  as  do  chlorine,  bromine,  and  iodine,  which  form  with 
it  the  corresponding  acids,  and  compounds  of  chlorine, 
bromine,  and  iodine  with  arsenic. 

Passed  into  a  solution  of  a  metallic  salt,  arsine  either 
reduces  the  salt  and  throws  down  the  metal  as  in  the 
case  of  silver  ;  or  it  forms  an  arsenide  of  the  metal,  acting 
in  this  case  very  much  as  hydrogen  sulphide  does  when 
passed  into  similar  solutions.  Considering  the  instability 
of  arsine,  it  is  not  surprising  that  it  acts  as  a  reducing 
agent.  It  will  be  remembered  that  hydriodic  acid  and 
hydrogen  sulphide  act  in  the  same  way  towards  some 
oxygen  compounds,  and  the  action  is  due  to  their  break- 
ing down  into  hydrogen  and  the  other  element.  Thus, 
when  hydriodic  acid  acts  as  a  reducing  agent  the  iodine 
is  left  uncombined,  and  when  hydrogen  sulphide  acts  in 
this  way  the  sulphur  is  left.  But  when  arsine  acts  as  a 
reducing  agent  both  the  hydrogen  and  the  arsenic  com- 


308  INORGANIC  CHEMISTRY. 

bine  with  oxygen.  Thus,  when  arsine  is  passed  into  a 
solution  of  silver  nitrate  this  reaction  take?  place  : 

AsH3  +  6AgN03  +  3H20  =  As(OH)3  +  6HNO3  +  6Ag. 

When,  on  the  other  hand,  arsine  is  passed  through  a 
solution  of  a  salt  of  a  difficultly  reducible  metal,  the  ar- 
senide of  the  metal  is  thrown  down  : 

2AsH3  +  3CuS04  =  As2Cu3  +  3H2SO4. 

Arsine  does  not  combine  with  acids  to  form  arsonium 
compounds  such  as  AsHJ,  analogous  to  ammonium  and 
phosphonium  compounds. 

There  is  a  second  compound  of  arsenic  and  hydrogen 
which  is  solid  and  appears  to  have  the  composition  rep- 
resented by  the  formula  As2H2. 

ANTIMONY,  Sb  (At.  Wt.  119.52). 

Occurrence. — Antimony  occurs  in  nature  chiefly  in  the 
form  of  stibnite,  which  is  the  trisulphide  Sb2S3.  This  also 
occurs  very  widely  distributed  in  nature  in  combination 
with  sulphides  of  various  metals,  as  copper,  lead,  and 
silver.  The  element  is  made  from  the  sulphide  either  by 
heating  it  with  iron,  with  which  the  sulphur  combines, 
leaving  the  antimony  free  ;  or  by  roasting  it,  that  is,  heat- 
ing it  in  combination  with  the  air,  thus  converting  the  anti- 
mony into  the  tetroxide  Sb2O4,  and  the  sulphur  into  the 
dioxide  SO2,  and  then  treating  the  oxide  of  antimony  with 
reducing  agents,  as,  for  example,  carbon : 

Sb204  +  4C  =  2Sb  +  4CO. 

Properties. — Antimony  is  hard  and  brittle  ;  has  a  silver- 
white  color ;  and  a  high  metallic  lustre.  It  can  be  dis- 
tilled at  white  heat.  At  ordinary  temperature  it  is  not 
changed  by  contact  with  the  air.  When  heated  to  a  suffi- 
ciently high  temperature  in  the  air  it  takes  fire  and  burns, 
forming  the  white  oxide  Sb2O3.  It  combines  directly  with 
chlorine,  forming  the  chloride  SbClB.  Nitric  acid  oxidizes 
it  either  to  antimony  oxide,  Sb2O3,  or  antimonic  acid, 


APPLICATIONS  OF  ANTIMONY—  STIBINE.          309 

H3SbO4.  Aqua  regia  dissolves  it.  Hot  concentrated  sul- 
phuric acid  dissolves  it,  forming  antimony  sulphate,  and 
sulphur  dioxide  escapes.  This  action  is  similar  to  that 
which  takes  place  when  sulphuric  acid  acts  upon  copper, 
It  is  probable  that  the  formation  of  the  sulphur  dioxide 
is  due  to  the  action  of  the  hydrogen  liberated  from  the 
sulphuric  acid  by  the  antimony  in  forming  antimony  sul- 
phate : 

2Sb        +  3H2S04  =  Sb2(S04)3  +  3H2  ; 
3H2SO4  +  3H2        =  3SO2        +  6H2O. 


This  power  to  replace  the  hydrogen  of  some  acids  dis- 
tinguishes antimony  from  arsenic  and  phosphorus,  while 
its  power  to  form  acids  corresponding  to  those  of  phos- 
phorus and  arsenic  shows  its  analogy  to  these  elements. 

Applications  of  Antimony.  —  Antimony  is  used  as  a 
constituent  of  several  alloys  which  are  somewhat  in- 
definite compounds  which  metallic  elements  form  with 
one  another.  Among  the  alloys  of  antimony  are  type- 
metal,  from  which  type  is  made,  and  britannia  metal. 
The  former  consists  of  lead  and  antimony,  and  the  latter 
of  tin  and  antimony.  There  are  a  number  of  alloys  which 
contain  antimony  which  will  be  referred  to  under  the 
other  constituents. 

Stibine,  SbH3.  —  This  analogue  of  ammonia,  phosphine, 
and  arsine  is  more  like  arsine  than  it  is  like  the  others. 
It  is  made  by  the  same  methods  as  those  used  in  making 
arsine,  i.e.,  by  treating  an  alloy  of  zinc  and  antimony 
with  sulphuric  acid,  or  by  reducing  oxides  of  antimony 
by  means  of  nascent  hydrogen.  The  latter  method  gives 
a  gas  which  contains  a  large  percentage  of  hydrogen,  but 
for  most  purposes  this  is  not  objectionable.  It  is  only 
necessary  to  introduce  into  a  flask  containing  zinc  and 
dilute  sulphuric  acid  a  little  of  a  solution  of  some  oxy- 
gen compound  of  antimony,  when  the  reduction  is  at  once 
effected,  and  the  escaping  hydrogen  contains  stibine. 

Stibine  is  a  colorless,  inodorous  gas,  which  burns  with 
a  greenish-white  flame.  In  general,  it  conducts  itself 
much  like  arsine.  It  is  unstable  and  breaks  down  when 
the  tube  through  which  it  is  passing  is  heated  to  red- 


310  INORGANIC  CHEMISTRY. 

ness.  It  then  leaves  a  deposit  which  looks  like  that 
formed  in  the  case  of  arsine.  When  a  cold  object,  as  a 
piece  of  porcelain,  is  held  for  a  moment  in  a  flame  of 
stibine  a  dark  deposit  is  formed  which  resembles  that 
formed  with  arsine. 

Methods  of  Distinguishing  between  Arsenic  and  Anti- 
mony.— As  arsenic  is  frequently  used  in  cases  of  poison- 
ing the  question  of  deciding  whether  it  is  present  in  a 
given  liquid  or  mixture  is  of  great  importance.  One  of 
the  chief  difficulties  encountered  is  the  similarity  of 
the  two  elements  arsenic  and  antimony.  The  method 
commonly  employed  in  examining  a  substance  for  arsenic 
is  known  as  Marsh's  test.  This  consists  in  getting  the 
substance  in  solution,  and  then  pouring  some  of  the 
liquid  into  a  vessel  containing  pure  zinc  and  pure  dilute 
sulphuric  acid.  If  arsenic  is  present  in  the  solution  it 
will,  under  these  circumstances,  be  converted  into  arsine, 
the  presence  of  which  can  be  recognized  by  heating  the 
tube  through  which  the  gas  is  passing,  and  by  holding  a 
piece  of  porcelain  in  the  flame.  If  deposits  are  not 
formed  in  the  tube  or  on  the  porcelain,  arsenic  is  not 
present ;  but  if  deposits  are  formed,  the  only  conclusion 
that  can  be  drawn  is  that  either  arsenic  or  antimony  is 
present,  or  possibly  both  may  be  present.  For  the  pur- 
pose of  distinguishing  between  the  two  elements,  advan- 
tage is  taken  of  the  following  differences  between  the 
spots  :  The  antimony  spots  are  darker  than  those  formed 
by  arsenic,  and  they  have  a  smoky  appearance,  while 
those  of  arsenic  have  not ;  further,  the  arsenic  deposits 
are  quite  volatile,  and  can  therefore  be  driven  before  the 
flame  in  the  tube  or  upon  the  porcelain,  while  those  of 
antimony  are  not  volatile ;  again,  the  deposits  of  arsenic 
are  easily  soluble  in  a  solution  of  sodium  hypochlorite 
or  hypobromite,  while  the  antimony  deposits  are  in- 
soluble in  these  solutions.  There  are  other  differences, 
but  those  mentioned  will  suffice  to  enable  a  careful 
worker  and  observer  to  distinguish  between  the  two  with- 
out any  possibility  of  doubt.  Another  difficulty  always 
encountered  in  examining  for  arsenic  is  the  fact  that  the 
sulphuric  acid,  the  zinc,  and  the  glass  of  which  the  ves- 


BISMUTH.  311 

sels  are  made  may  contain  arsenic.  It  is  quite  possible 
to  overcome  all  the  difficulties  and  to  decide  positively 
whether  arsenic  is  present  or  not.  If  it  is  found  that  on 
heating  the  tube  through  which  the  hydrogen  is  passing 
no  deposit  is  formed,  even  after  continued  heating,  and 
that  the  hydrogen  flame  gives  no  deposit  upon  a  piece 
of  porcelain  introduced  into  it,  then  it  is  safe  to  proceed 
with  the  examination  of  the  suspected  liquid.  If  the 
substance  which  is  to  be  examined  for  arsenic  has  to  be 
treated  with  chemical  compounds  in  order  to  prepare  it 
for  analysis,  every  compound  used  in  this  part  of  the 
process  must  be  separately  examined  for  arsenic. 

BISMUTH,  Bi  (At.  Wt.  206.54). 

Occurrence,  etc. — Bismuth  is  not  abundant  nor  widely 
distributed  in  nature.  It  occurs  for  the  most  part  native 
in  veins  of  granite  and  clay  slate.  Among  the  compounds 
of  bismuth  found  in  nature  are  the  oxide  Bi2O3  and  the 
corresponding  sulphide  Bi2S3. 

The  ores  are  roasted  and  then  treated  with  appropriate 
reducing  agents.  In  different  places  different  methods 
of  extraction  are  employed.  As  the  chief  applications  of 
bismuth  are  for  pharmaceutical  purposes,  it  is  necessary 
that  the  element  should  be  specially  pure ;  above  all, 
that  it  should  not  be  contaminated  with  arsenic.  In 
order  to  remove  the  last  traces  of  this  element  the  pow- 
dered bismuth  is  generally  melted  with  saltpeter. 

Bismuth  is  a  hard,  brittle,  reddish- white  substance 
with  a  metallic  lustre.  It  looks  very  much  like  antimony, 
but  is  distinguished  from  it  by  its  reddish  tint.  At  or- 
dinary temperatures  it  remains  unchanged  in  the  air. 
When  heated  to  red  heat  it  burns  with  a  bluish  flame, 
forming  the  yellow  oxide  Bi2O3. 

Hydrochloric  acid  scarcely  acts  upon  it ;  concentrated 
sulphuric  acid  forms  bismuth  sulphate,  Bi2(SO4)3,  in  which 
the  bismuth  evidently  plays  the  part  of  a  base-forming 
element ;  nitric  acid  gives  bismuth  nitrate,  Bi(NO3)3,  which 
is  partly  decomposed  by  water,  forming  so-called  basic 
nitrates  which  are  difficultly  soluble  in  water.  These 
salts  will  be  taken  up  in  the  next  chapter. 


312  INORGANIC  CHEMISTRY. 

Some  bismuth  is  used  in  the  preparation  of  alloys 
which  are  easily  fusible,  as,  for  example,  Newton's  metal, 
which  contains  bismuth,  lead,  and  tin ;  Hose's  metal, 
which  consists  of  the  same  constituents  in  slightly  dif- 
ferent proportions ;  and  Wood's  metal,  which  consists 
of  bismuth,  lead,  tin,  and  cadmium. 

Bismuth  does  not  combine  with  hydrogen. 

Compounds  of  the  Members  of  the  Phosphorus  Group 
with  the  Members  of  the  Chlorine  Group. — In  the  intro- 
duction to  this  chapter  it  was  stated  that  the  elements  of 
the  phosphorus  group  combine  with  chlorine  in  two 
proportions,  forming  compounds  of  the  general  formulas 
MC13  and  MC15.  Arsenic,  however,  forms  only  one  com- 
pound with  chlorine,  AsCl3,  while  bismuth  forms  one  of 
the  formula  BiCl3,  and  another,  Bi2Cl4.  The  compounds 
of  phosphorus  and  chlorine  are  the  best  known,  and  a 
brief  study  of  these  will  give  a  fair  idea  of  the  methods 
of  preparation  and  the  conduct  of  the  analogous  com- 
pounds of  the  other  members  of  the  group. 

Phosphorus  Trichloride,  PC13,  is  made  by  conducting 
dry  chlorine  gas  upon  phosphorus  in  a  retort  connected 
with  a  receiver.  Action  takes  place  at  once  with  evo- 
lution of  heat,  and  the  trichloride  distils  over  and  is 
condensed  as  a  liquid  into  the  receiver.  It  is  purified 
by  distillation  on  a  water-bath.  It  is  a  clear,  color- 
less liquid,  which  boils  at  74°.  In  contact  with  air  it 
fumes  in  consequence  of  the  action  of  the  water  vapor 
which  decomposes  it.  It  has  a  disagreeable  odor  of  its 
own  mixed  with  that  of  hydrochloric  acid.  Its  most 
characteristic  decomposition  is  that  which  it  undergoes 
with  water,  which  is  of  the  same  kind  as  that  which  the 
chlorides  of  sulphur,  selenium,  and  tellurium  undergo 
with  water.  The  general  tendency  of  the  chlorides  of 
the  acid-forming  elements  is  to  undergo  decomposition 
with  water  in  such  a  way  that  the  corresponding  hydroxyl 
compound  is  formed,  together  with  hydrochloric  acid. 
This  is  shown  in  the  case  of  tellurium  tetrachloride,  which 
with  water  forms  normal  tellurious  acid,  Te(OH)4,  and 
hydrochloric  acid  : 


PHOSPHORUS  TRICHLORIDE.  313 

HOH  f  OH 


HOH  [  OH 

In  the  case  of  sulphur  tetrachloride  the  hydroxyl  de- 
rivative, if  formed,  breaks  down  into  water  and  sulphur 
dioxide.  When  phosphorus  trichloride  is  treated  with 
water  the  decomposition  is  probably  represented  by  the 
equation 

HOH  (  OH 

PCI   4-  HOH  =  P^  OH  +  3HC1. 
HOH  (  OH 

From  some  experiments  it  appears  possible  that  this 
form  of  compound  is  unstable,  and  that,  owing  to  the 
marked  tendency  of  phosphorus  to  act  as  a  quinquivalent 
element,  the  constituents  arrange  themselves  differently, 

(H 

as  represented  in  the  formula   O=P-<  OH.    Thisques- 

(OH 

tion  will  be  referred  to  under  the  head  of  Phosphorous 
Acid. 

The  trichloride  shows  a  strong  tendency  to  take  up 
chlorine,  bromine,  iodine,  oxygen,  and  sulphur,  and  thus 
to  become  saturated  as  a  quinquivalent  element.  With 
chlorine  it  forms  the  pentachloride,  PC15,  with  oxygen 
the  oxychloride,  POC13,  and  with  sulphur  the  sulphochlo- 
ride,  PSC13.  It  does  not,  however,  readily  take  up  free 
oxygen  or  free  sulphur  directly,  but  will  take  up  these 
elements  from  compounds  in  which  they  are  not  firmly 
held.  Thus,  when  the  trichloride  is  brought  together 
with  sulphur  trioxide  this  reaction  takes  place  : 

S03  +  PC13  =  POC13  +  SO2  ; 

and  when  it  is  brought  together  with  a  polysulphide, 
as  Na2S&,  it  takes  up  a  part  of  the  sulphur  and  forms 
the  sulphochloride,  PSC13.  So,  further,  it  is  converted 
into  the  oxychloride  when  treated  with  ozone.  These 
reactions  show  the  marked  tendency  which  the  trichlo- 


314  INORGANIC  CHEMISTRY. 

ride  lias  to  pass  over  into  compounds  of  quinquivalent 
phosphorus — a  tendency  which  is  characteristic  of  phos- 
phorus compounds  in  general. 

Phosphorus  Pentachloride,  PC15,  is  formed  by  treating 
phosphorus  or  the  trichloride  with  dry  chlorine.  It 
is  best  prepared  by  passing  chlorine  through  a  wide 
tube  upon  the  surface  of  the  trichloride,  contained 
in  a  vessel,  which  is  kept  cool.  Gradually  the 
liquid  becomes  thicker  and  thicker,  and  finally,  if  well 
stirred,  it  becomes  solid.  It  is  a  white  solid,  but  it 
generally  has  a  slightly  yellowish  or  greenish  color  in 
consequence  of  a  slight  decomposition  into  the  tri- 
chloride and  free  chlorine.  It  sublimes  below  100° 
without  melting.  When  heated  to  boiling  it  under- 
goes partial  decomposition  into  chlorine  and  the  trichlo- 
ride, and  this  decomposition  is  complete  at  about  300°. 
As  the  temperature  is  raised  from  the  apparent  boiling 
point  to  the  point  at  which  the  decomposition  is  com- 
plete, the  color  of  the  vapor  is  seen  to  grow  darker 
in  consequence  of  the  increased  quantity  of  free  chlo- 
rine present.  The  decomposition  is  gradual,  and,  for 
any  given  temperature,  the  amount  of  decomposition 
is  constant.  This  kind  of  decomposition,  which  is  known 
as  dissociation,  has  been  studied  very  carefully,  and  is 
found  to  be  capable  of  explanation  by  the  aid  of  the 
kinetic  theory  of  gases.  In  a  later  chapter  this  subject 
will  be  treated,  and  a  number  of  other  examples  will  be 
given.  Owing  to  this  decomposition  under  the  influence 
of  heat  the  specific  gravity  of  the  vapor  of  phosphorus 
pentachloride  is  not  what  it  should  be,  if  the  formula  is 
PC15.  On  the  other  hand,  the  specific  gravity  of  the 
vapor  of  the  trichloride  leads  to  the  formula  PC13,  and 
that  of  the  oxychloride  to  the  formula  POC13.  The  ap- 
parent anomaly  presented  by  the  pentachloride  is  easily 
understood.  When  a  molecule  of  the  compound  is  con- 
verted into  vapor,  or  is  heated  to  a  sufficiently  high 
temperature,  it  is  broken  down  in  accordance  with  this 
equation : 

PC16  =  PC13  +  C12. 


PHOSPHORUS  PENTACHLOEIDE.  315 

From  the  one  molecule,  therefore,  two  gaseous  mole- 
cules are  'obtained.  Consequently  the  vapor  formed  oc- 
cupies twice  as  much  space  as  it  would  if  there  were  no 
decomposition.  It  follows  that  the  specific  gravity  of 
the  vapor  must  be  only  half  what  it  would  be  if  there 
were  no  decomposition.  When  the  compound  is  con- 
verted into  vapor  in  an  atmosphere  of  phosphorus  tri- 
chloride, the  decomposition  referred  to  does  not  take 
place,  and,  under  these  circumstances,  the  specific  gravity 
is  found  to  be  in  accordance  with  Avogadro's  law,  and 
with  the  formula  PC15.  This  case  is  a  particularly  in- 
teresting one,  as  it  has  played  an  important  part  in 
the  discussions  in  regard  to  the  validity  of  Avogadro's 
law. 

The  conduct  of  phosphorus  pentachloride  towards 
water  is  in  general  like  that  of  the  other  chlorides  of 
acid-forming  elements.  But,  owing  probably  to  a  second- 
ary action,  the  product  is  not  the  corresponding  hydroxyl 
compound.  It  is  probable  that  the  first  action  of  the 
water  is  represented  by  the  equation 

TTTT/-V  (    OH 

PC1       *~      =  P    OH      2H01. 


But  this  product,  if  formed,  breaks  down  at  once  into 
phosphorus  oxychloride  and  water,  and  the  water  thus 
given  off  acts  upon  a  further  quantity  of  the  penta- 
chloride : 


The  formation  of  the  oxychloride  from  the  penta- 
chloride by  the  action  of  water  takes  place  very  easily. 
The  oxychloride  is  then  further  acted  upon  by  the  water, 
and  each  chlorine  atom  is  replaced  by  hydroxyl : 

HOH  ( OH 

OPC13  +  HOH  =  OP^  OH  +  3HC1. 
HOH  OH 


316  INORGANIC  CHEMISTRY. 

The   final  product  is   the   acid   H3PO4,  or  phosphoric 
acid. 

It  will  be  seen  that  the  effect  of  phosphorus  penta- 
chloride  upon  water  is  to  replace  the  hydroxyl  of  the 
water  by  chlorine.  Thus,  one  molecule  of  the  penta- 
chloride  and  five  molecules  of  water  give  one  molecule 
of  phosphoric  acid  and  five  of  hydrochloric  acid  : 

HOH  HC1 

HOH  HC1 

PC15  +  HOH  =  OP(OH)3  +  HC1  +  H2O. 
HOH  HC1 

HOH  HC1 

In  the  reaction,  the  hydroxyl  of  the  water  and  the 
chlorine  of  the  chloride  exchange  places.  Similarly, 
when  any  compound  which  contains  hydroxyl  is  treated 
with  phosphorus  pentachloride  the  same  reaction  takes 
place,  the  hydroxyl  being  replaced  by  chlorine.  There- 
fore phosphorus  pentachloride  may  be  used  as  a  reagent 
for  testing  for  the  hydroxyl  condition  in  compounds.  If 
a  compound  which  contains  hydrogen  and  oxygen  is 
treated  with  the  pentachloride,  and  an  atom  of  hydrogen 
and  one  of  oxygen  is  replaced  by  an  atom  of  chlorine, 
the  conclusion  is  drawn  that  the  compound  contains 
hydroxyl.  This,  of  course,  amounts  to  saying  that  the 
compound  resembles  water  in  its  reaction  with  the  penta- 
chloride, and  this  is  most  easily  explained  by  the  as- 
sumption that  the  same  condition  exists  in  both.  It 
should  be  borne  in  mind,  further,  that,  in  general,  any 
compound  of  chlorine  with  an  acid-forming  element 
which  undergoes  decomposition  with  water  might  be  used 
for  the  same  purpose.  The  action  of  the  pentachloride 
upon  a  hydroxyl  compound  is  well  illustrated  in  the  case 
of  sulphuric  acid : 

SO'<OH  +  PC1*  =  POC1*  +  SO«<OH  +  HCL 

The  action  of  the  oxychloride  would  take  place  as  rep- 
resented thus : 


ARSENIC  TRICHLORIDE.  31 

POC1°  = 


Phosphorus  forms  a  tribromide,  PBr3  ,  and  a  penta- 
bromide,  PBrB,  similar  to  the  chlorides.  With  iodine  it 
forms  the  compounds  P,I4  and  PI3,  and  with  fluorine  a 
trifluoride,  PF3,  and  a  pentafluoride,  PFB. 

Arsenic  Trichloride,  AsCl3,  is  the  only  compound  which 
arsenic  forms  with  chlorine.  It  is  easily  made  by  passing 
chlorine  into  a  retort  containing  powdered  arsenic.  It 
is  also  formed  by  the  action  of  phosphorus  pentachloride 
on  arsenic  trioxide,  in  which  case  the  oxygen  of  the  tri- 
oxide  is  simply  replaced  by  chlorine  : 

As203  +  3PC15  =  3POC13  +  2AsCl3. 

Further,  it  is  obtained  when  dry  hydrochloric  acid  gas 
is  passed  over  arsenic  trioxide,  and  it  is  present  in  solu- 
tion when  the  trioxide  is  dissolved  in  strong  hydrochloric 
acid. 

Arsenic  trichloride  is  a  liquid  which  boils  at  about 
134°.  It  is  extremely  poisonous.  If  a  hydrochloric  acid 
solution  of  the  trichloride  is  boiled,  it  passes  over  when 
the  temperature  rises  above  100°.  When  substances 
containing  arsenic  are  heated  in  a  retort  with  sulphuric 
acid  and  sodium  chloride,  arsenic  chloride  is  formed  and 
passes  over  into  the  receiver. 

When  arsenic  trichloride  is  treated  with  a  little  water 
two  of  the  chlorine  atoms  are  replaced  by  hydroxyl,  thus  : 


2HC1. 


If  treated  with  much  water  it  yields  arsenious  acid  ;_  .  . 
AsCl3  +  3H2O  =  As(OH)3  +  3HCL 

We  have  here  a  curious  illustration  of  the  effect  of  the 
relative  quantities  of  the  substances  which  take  part  in 
a  reaction  upon  the  nature  of  the  reaction.  When 


318  INORGANIC  CHEMISTRY. 

arsenic  trioxide  is  treated  with  an  excess  of  hydrochloric 
acid  it  is  converted  into  the  chloride : 

As2O3  +  6HC1  =  2AsCl3  +  3H2O. 

If,  on  the  other  hand,  the  chloride  is  treated  with  a 
large  excess  of  water  it  is  completely  broken  down,  yield- 
ing the  trihydroxyl  derivative,  or  arsenious  acid.  There 
are  many  other  illustrations  of  this  kind  met  with  among 
chemical  reactions,  and  the  subject  of  the  influence  of 
mass  in  determining  the  character  of  a  reaction  has  been 
studied  with  much  care.  This  subject  will  receive  at- 
tention in  a  later  chapter  under  the  head  of  Mass  Action. 
Compounds  of  Antimony  and  Chlorine. — Antimony  forms 
two  compounds  with  chlorine,  analogous  in  compo- 
sition to  the  two  chlorides  of  phosphorus.  These 
are  antimony  trichloride,  SbCl8,  and  antimony  penta- 
chloride,  SbCl6.  The  trichloride  is  formed  by  direct 
treatment  of  the  element  with  chlorine.  It  is  also 
formed  by  dissolving  antimony  in  hydrochloric  acid 
with  addition  of  nitric  acid,  and  when  this  solution 
is  distilled  the  chloride  of  antimony  passes  over.  At  the 
ordinary  temperatures  it  is  a  solid,  colorless,  crystal- 
line, soft  substance,  which,  on  account  of  its  consistency, 
has  received  the  common  name  "  butter  of  antimony  " 
(Butyrum  Antimonii).  When  treated  with  water  it  forms 
an  oxychloride,  the  composition  of  which  varies  according 
to  circumstances,  but  it  generally  approximates  to  that 
represented  by  the  formula  Sb4O6Cl2.  If  treated  with 
cold  water  the  decomposition  is  simple,  the  product 
being  antimony  oxychloride,  SbOCl : 

SbCl3  +  H20  =  SbOCl  +  2HC1. 

The  oxychlorides  formed  by  treating  antimony  tri- 
chloride with  hot  water,  which,  as  already  stated,  have 
a  composition  approximating  that  represented  by  the 
formula  Sb4O6Cl2,  are  known  as  "  Powder  of  Algaroth" 
The  relation  of  this  compound  to  the  simple  oxychloride, 
SbOCl,  is  indicated  by  the  equation 


DOUBLE  SALTS.  319 

SSbOCl  =  Sb406Cla  +  SbCl3. 

Liquor  Stibii  Muriatici  is  a  solution  of  antimony  tri- 
chloride prepared  by  dissolving  antimony  sulphide  in 
concentrated  hydrochloric  acid : 

Sb2S3  +  6HC1  =  2SbCl3  +  3H2S. 

Antimony  chloride  has  a  caustic  action  and  is  used  in 
medicine.  It  is  also  used  for  the  purpose  of  burnishing 
iron  ware,  as  gun-barrels.  When  treated  with  the  chloride 
they  acquire  a  brownish,  bronze-like  color. 

PentachLoride  of  Antimony  is  in  general  like  the  pen- 
tachloride  of  phosphorus.  It  gives  up  its  chlorine, 
however,  somewhat  more  readily.  When  treated  with 
water  it  yields  first  an  oxychloride,  SbO2Cl,  and  this  is 
further  acted  upon  and  converted  into  antimonic  acid, 
H3Sb04. 

Bismuth  and  Chlorine. — Bismuth  differs  from  the  other 
members  of  the  phosphorus  group  in  its  conduct  towards 
chlorine.  It  forms  the  chloride,  Bi2Cl4,  of  which 
there  is  no  analogue  among  the  compounds  of  the  other 
elements  of  the  group  with  chlorine.  The  compound 
may,  however,  be  regarded  as  analogous  to  the  hydrogen 
compounds  hydrazine,  N2H4,  and  liquid  phosphine,  P2H4. 
Bismuth  dichloride,  Bi2Cl4,  is  formed  by  reduction  of  the 
trichloride,  BiCl3,  when  the  latter  is  treated  with  hydro- 
gen at  a  temperature  of  about  300°.  Bismuth  trichloride, 
BiCl3,  is  formed  by  treating  bismuth  with  chlorine,  and 
by  dissolving  bismuth  oxide,  BiQO3,  in  hydrochloric  acid. 
From  this  solution  a  crystallized  compound  of  the 
formula  BiCl3  +  H2O  is  deposited,  and  it  is  impossible 
to  drive  all  the  water  off  from  this  compound  without 
causing  decomposition.  Treated  with  water  the  chloride 
is  decomposed,  forming  the  oxychloride,  BiOCl : 

BiCl3  +  H2O  =  BiOCl  +  2HC1. 

Double  Salts. — When  the  chlorides,  bromides,  iodides, 
and  fluorides  of  the  members  of  the  phosphorus  group  are 
treated  with  the  corresponding  salts  of  potassium,  sodium, 


320  INORGANIC  CHEMISTRY. 

and  some  other  strongly  marked  base-forming  elements, 
compounds  known  as  double  salts  are  formed.  Thus, 
when  antimony  trichloride  is  treated  with  a  strong  water 
solution  of  potassium  chloride,  a  salt  of  the  composition 
SbCl3.3KCl  is  formed.  This  compound  appears  to  be 
analogous  to  the  oxygen  compound  K3SbO3,  or  potassium 
antimonite,  differing  from  it  in  containing  chlorine  in 
place  of  the  oxygen.  While  the  latter  is  derived  from 
an  acid  of  the  composition  H3SbO3,  the  former  appears 
to  be  derived  from  the  corresponding  chlorine  acid, 
H3SbCl6.  Similar  compounds  of  sulphur  and  tellurium 
are  known,  and  one  was  referred  to  on  page  205.  This 
is  the  potassium  compound  K2TeBr6  or  TeBr4.2KBr, 
which,  as  explained  there,  is  analogous  to  potassium 
tellurite,  K3Te03. 


CHAPTER  XVIII. 

COMPOUNDS  OF  THE  ELEMENTS  OF  THE  PHOSPHORUS 
GROUP  WITH  OXYGEN  AND  WITH  OXYGEN  AND  HY- 
DROGEN. 

Introduction. — The  product  of  the  direct  action  of 
oxygen  upon  phosphorus  is  the  pentoxide  PA-  Ar- 
senic, antimony,  and  bismuth,  however,  form  the  tri- 
oxides  As2O3,  Sb2O3,  and  Bi2O3.  It  is  possible  to  obtain 
a  compound  of  arsenic  and  oxygen  of  the  formula  As2O6, 
one  of  antimony,  Sb2O4,  and  another,  Sb  A,  and,  finally, 
two  oxides  of  bismuth,  Bi2O2  and  Bi2O5.  Phosphorus, 
further,  forms  the  oxides  P4O,  P,O8,  and  P3O4.  The 
table  below  contains  the  formulas  of  the  above-men- 
tioned compounds  systematically  arranged : 

P40 


PA         AsA         Sb,0, 

PA  Sb,o4 

PA        AsA        SbA 

The  final  products  of  the  oxidation  of  the  elements  of 
this  group,  if  water  is  present,  are  phosphoric,  arsenic, 
antimonic,  and  bismuthic  acids.  All  of  these  are  well- 
marked  acids  except  the  last.  They  can  all  be  regarded 
as  derived  from  the  normal  acids  of  the  general  formula 
M(OH)5  by  loss  of  one  or  two  molecules  of  water.  The 
common  forms  of  phosphoric,  arsenic,  and  antimonic 
acids  are  those  which  are  formed  from  the  normal  acids 
by  loss  of  one  molecule  of  water : 

P(OH)6  P{(OH)S        +         H»0; 

Normal  phosphoric  acid  Orthophosphoric  acid 


As(OH)5  As{( 


(OH)s  > 

Normal  arsenic  acid  Orthoarsenic  acid 

(321) 


322  INORQANIG  CHEMISTRY. 

Sb(OH)s  Sb{(OH)s 

Normal  antimonic  acid  Orthoantimonic  acid 


Bismuthic  acid  appears,  however,  to  be  formed  from 
the  normal  acid  by  loss  of  two  molecules  of  water,  just 
as  the  so-called  metaphosphoric,  metarsenic,  and  metanti- 
monic  acids  are : 

Bi(OH)5        =          BiJQH  +       2H20; 

Normal  bismuthic  acid  Bismuthic  acid 


P(OH)5  P|gk  +        2H,0; 

Metaphosphoric  acid 

As(OH)5        =         As  -j  Q|J  -f-        2H2O ; 

Metarsenic  acid 

Sb(OH)5       =         sM8b  +        2H'°- 

Metantimonic  acid 

From  the  ordinary  or  ortho  acids,  and  from  the  meta 
acids,  more  complex  forms  can  be  derived  by  loss  of 
different  quantities  of  water.  The  most  common  form 
besides  those  mentioned  is  that  seen  in  the  so-called 
pyro  acids,  of  which  pyrophosphoric  acid  is  the  best 
known  example.  It  is  formed  from  the  ortho  acid  by 
loss  of  one  molecule  of  water  from  two  molecules  of 
the  acid,  just  as  pyrosulphuric  or  disulphuric  acid  is 
formed  from  two  molecules  of  ordinary  sulphuric  acid 
by  loss  of  one  molecule  of  water.  The  formation  of 
pyrophosphoric  acid  from  orthophosphoric  acid  takes 
place  according  to  the  equation 


OH 
OH 

OH 
pi  OH  p " 

18H  A  ; 


COMPOUNDS  OF  THE  PHOSPHORUS  GROUP.        323 

or 

2H.PO,        =        HJP.O,        +        H,0. 

Orthophosphoric  acid  Pyrophosphoric  acid 

Pyroarsenic  and  pyroantimonic  acids  bear  the  same 
relations  to  the  ortho  acids  that  pyrophosphoric  acid 
bears  to  orthophosphoric  acid. 

By  partial  oxidation  of  phosphorus  in  presence  of 
water,  phosphorous  acid,  H3PO3,  is  formed.  The  same 
acid  is  formed  by  the  action  of  phosphorus  trichloride 
on  water.  According  to  the  latter  method  of  formation 
we  should  expect  to  find  that  this  acid  is  normal  phos- 
phorous acid,  P(OH)3.  As  already  stated,  however,  it 
appears  probable  that  the  acid  has  the  constitution 

/H 
O=P—  -OH.      The   acids   of   arsenic   and    antimony   of 

\OH 

similar  composition  seem  to  be  the  normal  acids  As(OH)8 
and '  Sb(OH)3.  The  hydroxyl  derivative  of  bismuth 
corresponding  to  these  acids  has  no  acid  properties, 
but  on  the  contrary  is  basic.  Hypophosphorous  acid 
has  the  composition  H3PO2.  It  is  monobasic,  and  it 
appears  therefore  that  it  contains  but  one  hydroxyl,  as 
represented  in  the  formula  H2OP(OH).  It  is  possible 
that  the  relations  between  phosphoric,  phosphorous, 
and  hypophosphorous  acids  should  be  represented  by 
the  formulas 

(OH  (  H  (  H 

OP^OH,        OP^OH,        OP^H  . 
(OH  (OH  (OH 

Phosphoric  acid  Phosphorous  acid  Hypophosphorous  acid 

The  fundamental  compound,  then,  from  which  these  may 
be  regarded  as  derived  is  the  unknown  oxyphosphine 
OPH3.  By  oxidation  we  should  expect  phosphine 
to  yield  in  successive  stages  the  three  products  above 
named : 

(H  (H  (H  (H  (OH 

P^H,OP^H,     OP^H    ,     OP^OH,OP^OH. 

(H        (H         (OH         (OH        (OH 

Unknown     Hypophosphorous       Phosphorous       Phosphoric 
acid  acid  acid 


3M  INORGANIC  CHEMISTRY. 

The  oxidation  of   hydrogen  sulphide  takes  place  sim* 
ilarly,  as  has  been  shown  : 


,       , 

Unknown  Sulphurous  acid  Sulphuric  acid 

With  oxygen  and  chlorine  the  elements  of  the  phos- 
phorus group  form  a  number  of  compounds  known  as 
oxychlorides.  Towards  chlorine  as  well  as  towards 
oxygen  all  these  elements  except  bismuth  are  quin- 
quivalent. A  part  or  all  of  the  oxygen  of  the  oxygen 
compounds  can  be  replaced  by  chlorine.  Starting  with 
the  chlorine  compound  on  the  one  hand,  oxychlorides. 
can  be  obtained  from  it,  until  all  the  chlorine  is  replaced 
by  oxygen,  and  the  limit  is  reached  in  the  oxide.  So 
also  the  chlorine  can  be  replaced  by  hydroxyl  and  the 
acids  thus  obtained. 

(1)  PC15  gives  POC13  and  P2O6  as  final  product  ; 

(2)  PC15  gives  POC13  and  with  water  PO(OH)3. 

(3)  PC13  gives  as  final  product  P2O3  ; 

(4)  PC13  gives  with  water  P(OH)3. 

Intermediate  products  are  supposable,  but  not  known, 
as,  for  example  : 


Cl  (Cl  (Cl  (OH 

Cl         P^  Cl  ,        PK  OH,       Pi  OH. 
Cl  OH  OH  OH 


(  Cl 
A  compound  of  arsenic   of  the  formula  As-j  /QTT\     is 

known,  however,  and  this  plainly  corresponds  to  one  of 
these  intermediate  products. 

With  sulphur  phosphorus  apparently  forms  a  large 
number  of  compounds.  Among  them  are  two  which 
have  the  formulas  P2SS  and  P2S5.  These  plainly  are 
analogous  to  the  two  oxides  of  phosphorus,  P2O3  and 
P2OB.  When  treated  with  water  these  sulphur  com- 
pounds like  the  corresponding  chlorine  compounds 
yield  the  oxygen  acids.  Thus  the  trisulphide  undergoes 


COMPOUNDS  OF  THE  PHOSPHORUS  GROUP.        325 

decomposition  with  water  according  to  the  following 
equation : 

P2S3  +  6H2O  =  2H3PO3  +  3HaS ; 

and  the  pentasulphide  is  converted  by  water  into  phos* 
phoric  acid : 

PaS6  +  8HaO  =  2H3P04  +  5HaS. 

Arsenic  forms  with  sulphur  several  compounds,  the 
principal  of  which  are  the  disulphide,  As2S2,  the  trisul- 
phide,  As2S3,  and  the  pentasulphide,  As2S6.  The  principal 
sulphides  of  antimony  are  those  of  the  formulas  Sb2S9 
and  Sb2S5,  and  of  bismuth  those  of  the  formulas  Bi2Sa 
and  Bi2S3.  In  general,  therefore,  the  sulphur  compounds 
are  analogous  in  composition  to  the  oxygen  compounds, 
while  the  number  of  sulphur  compounds  of  these  ele- 
ments is  larger  than  that  of  the  oxygen  compounds. 
The  formulas  of  the  principal  sulphur  compounds  of  this 
group  are  given  systematically  arranged  in  the  table 
below : 

As2Sa        Bi2Sa 

P2S3        As2S3        SbaS8        BiaS3 
PaS5        As2S,        SbaS6        

Further,  there  are  sulphur  acids  which  are  to  be  re- 
garded as  the  oxygen  acids,  a  part  or  all  of  whose 
oxygen  is  replaced  by  sulphur.  Thus,  in  the  case  of 
arsenic  the  following  possibilities  suggest  themselves, 
starting  with  arsenious  acid : 

(SH 

-{  SH ; 

(SH 

and  starting  with  arsenic  acid,  the  following  possibilities 
suggest  themselves : 

OH  (OH  (OH  (OH          ( 

OH ,  SAs^  OH ,  SAs^  OH ,  SAs^  SH,  SAsJ 

OH        (OH        (SH        (SH        (SH 


326  INORGANIC  CHEMISTRY. 

While  none  of  these  compounds  is  known,  many  com- 
pounds are  known  which  are  to  be  regarded  as  salts  of 
one  or  another  of  these  acids.  Thus  salts  of  the  general 
formulas  M3AsS3  and  M3AsS4  are  well  known,  as  are  also 
salts  of  the  general  formula  MAsS2,  which  are  derived 

(  ^ 
from   the   acid  As  -j  QTT  ,  corresponding   to  the  oxygen 

compound  As  j  QTT  ,  which  in  turn  is  derived  from  arsen- 
ious  acid  by  loss  of  one  molecule  of  water  : 


So,  too,  we  have  : 

SH  , 

+H2S. 


Similar  compounds  of  antimony  are  also  well  known. 
The  possibility  of  making  analogous  compounds  contain- 
ing selenium  and  tellurium  will  suggest  itself. 

Phosphoric  Acid,  Orthophosphoric  Acid,  H3PO4.  —  The 
compound  to  which  the  name  phosphoric  acid  is  gener- 
ally applied,  and  from  which  the  best  known  phosphates 
are  derived,  is  that  which  has  the  formula  H3PO4.  To 
distinguish  it  from  the  other  varieties  it  is  called  ortho- 
phosphoric  acid.  As  has  been  stated,  this  is  the  final 
product  of  oxidation  of  phosphorus  in  the  presence  of 
water.  Thus,  when  phosphorus  is  boiled  with  nitric  acid 
it  is  converted  into  orthophosphoric  acid  ;  and  also  when 
phosphorus  is  burned  in  the  air,  and  the  product  dis- 
solved in  water,  phosphoric  acid  is  formed.  In  this  case 
the  first  product  of  the  oxidation  is  the  pentoxide  P2O5, 
also  known  as  phosphoric  anhydride,  and  when  this  is 
treated  with  water  it  is  converted  into  phosphoric  acid  : 

Pa06  +  3H3O  =  2H3PO4. 
The  occurrence  of  phosphoric  acid  in  nature  has  already 


PHOSPHORIC  ACID.  327 

been  referred  to  in  connection  with  the  occurrence  of 
phosphorus,  which  is  found  in  nature  almost  exclusively 
in  the  form  of  phosphates,  principally  as  calcium  phos- 
phate, Ca3(PO4)2,  in  phosphorite,  apatite,  and  the  ashes 
of  bones.  It  is  formed  when  either  phosphorus  penta- 
chloride  or  the  oxychloride  is  decomposed  by  water  : 

PC15    +  4H20  =  PO(OH)3  +  5HC1  ; 
POC13  +  3H2O  =  PO(OH)3  +  3HC1  ; 

and  from  the  analogous  bromine  and  iodine  compounds 
in  the  same  way.  In  order  to  prepare  the  acid  two  ways 
suggest  themselves  :  (1)  by  oxidizing  phosphorus  with 
nitric  acid  ;  and  (2)  by  extracting  it  from  one  of  the  natu- 
ral phosphates,  as  phosphorite  or  bone-ash.  The  first 
of  these  methods  is  better  adapted  to  the  preparation  of 
pure  phosphoric  acid,  such  as  is  needed  for  medicinal 
purposes  ;  the  latter  is  used  where  absolute  purity  of  the 
product  is  not  required.  It  should  be  said,  however, 
that  the  acid  obtained  by  oxidation  of  phosphorus  is  not 
pure,  as  commercial  phosphorus  almost  always  contains 
arsenic  and  small  quantities  of  other  impurities.  The 
arsenic  can  easily  be  removed  by  passing  hydrogen  sul- 
phide through  the  solution  after  the  nitric  acid  has  been 
evaporated.  If  the  solution  is  then  filtered  and  evapo- 
rated to  dryness,  the  orthophosphoric  acid  is  transformed 
into  pyrophosphoric  or  metaphosphoric  acid  according 
to  the  temperature  : 


H3P04  =  HP03   +H20. 

The  preparation  of  phosphoric  acid  from  a  phosphate 
is  not  a  simple  matter.  If  the  acid  were  volatile  or  in 
soluble  there  would  be  no  difficulty  in  separating  it.  In 
the  former  case  it  would  only  be  necessary  to  proceed  as 
in  preparing  hydrochloric  and  nitric  acids.  By  adding 
an  acid  which  is  not  volatile  except  at  a  high  temperature, 
such,  for  example,  as  sulphuric  acid,  and  heating,  the 
non-volatile  acid  replaces  the  volatile.  On  the  other 


328  INORGANIC  CHEMISTRY. 

hand,  if  phosphoric  acid  were  insoluble  in  water,  it  could 
be  separated  by  adding  a  soluble  acid  to  one  of  its  soluble 
salts.  When,  for  example,  nitric  acid  is  added  to  a  solu- 
tion of  potassium  tellurite,  KaTeO3,  tellurious  acid,  being 
insoluble,  is  thrown  down  : 

K2Te03  +  2HN03  =  2KNO3  +  HaTeO3. 

But  phosphoric  acid  is  not  volatile  and  is  soluble,  so  that 
plainly  neither  of  these  methods  can  be  used.  By  treat- 
ing the  calcium  salt  with  sulphuric  acid  the  calcium  can 
be  completely  separated  in  the  form  of  calcium  sul- 
phate, which  is  difficultly  soluble  in  water  and  insoluble 
in  alcohol.  The  ideal  reaction  to  be  accomplished  is  that 
represented  in  the  following  equation  : 

Ca3(P04)2  +  3H2S04  =  3CaS04  +  2H3P04. 

But  when  sulphuric  acid  is  added  to  calcium  phosphate, 
only  a  part  of  the  calcium  is  thrown  down  as  sulphate, 
the  rest  remaining  in  the  form  of  primary  calcium  phos- 
phate : 

Ca/PO,),  +  2H,SO,  =  2CaSO.  +  CaH,(PO4)s. 


The  phosphate  thus  formed  is  soluble  in  water,  and  the 
calcium  is  not  easily  precipitated  from  it.  By  evapora- 
tion and  addition  of  sufficient  sulphuric  acid  and  alcohol 
the  precipitation  can  be  effected,  and  a  solution  of  phos- 
phoric acid  thus  obtained.  This  acid  is  not  pure,  as 
there  are  substances  in  bone-ash  which  are  not  removed 
by  the  method  described. 

Phosphoric  acid  for  medicinal  purposes  is  often  made 
by  dissolving  the  pentoxide  in  water. 

Properties.  —  When  evaporated  to  the  proper  consis- 
tency the  acid  forms  a  thick  syrup  which  slowly  solidifies 
in  the  form  of  large  crystals.  The  crystals  are  deliques- 
cent. When  heated  to  a  sufficiently  high  temperature 
the  acid  loses  water,  as  already  explained,  and  yields, 
first,  pyrophosphoric,  and  then  metaphosphoric  acid. 
It  is  a  tribasic  acid,  capable  of  yielding  three  classes  of 


PHOSPHATES.  329 

(OH          (OH 

salts  of  the  general  formulas  OP-<  OH,  OP-<  OM,and 

(  OM          ( OM 

(OM 

OP  -<  OM ,  which  are  known  respectively  as  the  primary, 

(OM 

secondary,  and  tertiary  phosphates.  The  primary  and 
secondary  phosphates  are  also  known  as  acid  phosphates, 
and  the  tertiary  salts  as  neutral  or  normal  phosphates.  In 
these  salts  it  is  not  necessary  that  all  the  hydrogen  should 
be  replaced  by  the  same  metal.  There  are  salts  in  which 
two  or  three  metals  take  the  place  of  the  hydrogen  atoms. 
A  phosphate  much  used  in  the  laboratory,  for  example, 
is  one  in  which  one  hydrogen  atom  of  phosphoric  acid  is 
replaced  by  a  sodium  atom,  and  another  by  the  ammoni- 

(OH 
um  group,  NH,.      This  salt  has  the  formula  OP-<  ONa  , 

(ONH. 

and  is  called  ammonium  sodium  phosphate.  Another 
phosphate  commonly  met  with  is  ammonium  magnesium 

(ONH4 
phosphate,  OP-<  O    jyr   ,  which  is  derived  from  the  acid 

by  replacement  of  two  hydrogen  atoms  in  the  molecule 
by  one  bivalent  magnesium  atom,  and  one  by  the  am- 
monium group.  The  changes  which  these  three  classes 
of  phosphates  undergo  when  heated  are  of  special  inter- 
est. The  tertiary  phosphates  are  stable.  The  primary 
and  secondary  phosphates  give  up  all  their  hydrogen, 
which  passes  off  in  the  form  of  water.  Thus,  primary 
sodium  phosphate,  H2NaPO4,  loses  one  molecule  of  water 
from  each  molecule  of  the  salt,  and  is  converted  into  the 
metaphosphate,  NaPO3 : 

(OH 

OP^  OH    =  O2P(ONa)  +  H,0. 
(ONa 

In  general,  the  primary  phosphates  are  converted  into  meta- 
phosphates  by  heat. 

When  a  secondary  phosphate  is  heated  the  product  is 


330  INORGANIC  CHEMISTRY. 

a  pyrophosphate,  as  when  secondary  sodium  phosphate 
is  heated  to  a  sufficiently  high  temperature  it  is  converted 
into  sodium  pyrophosphate  : 


w  >- 

j  ONa         OP    ONa 

-<  ui>a  J  Q^_ 

(ONa 

In  general,  a  secondary  phosphate  is  converted  into  a  py- 
rophosphate by  heat. 

The  above  rules  do  not  hold  good  for  ammonium  salts, 
for  these  always  undergo  another  kind  of  decomposition 
when  heated.  When  sodium  ammonium  phosphate  is 
heated,  ammonia  is  first  given  off,  thus  : 

HNa(NH4)P04  =  H2NaPO4  +  NH3  ; 

and  the  primary  salt  formed  breaks  down  according  to 
the  above  rule,  forming  the  metaphosphate.  So,  also, 
when  ammonium  magnesium  phosphate  is  heated,  the 
first  change  consists  in  the  giving  off  of  ammonia,  thus  . 

(NH,)MgPO,  =  HMgPO,  +  NHS  ; 

and  the  secondary  magnesium  phosphate  thus  formed 
then  breaks  down,  forming  the  pyrophosphate,  Mg2P2O7  : 


The  presence  of  phosphoric  acid  can  be  detected  by 
means  of  the  following  characteristic  reactions  :  With 
silver  nitrate  it  gives  a  yellow  precipitate  of  tertiary  sil- 
ver phosphate,  Ag3PO4  ;  with  a  soluble  magnesium  salt 
and  ammonia  it  gives  ammonium  magnesium  phosphate, 
(NH4)MgPO4,  which  is  insoluble  in  water  ;  with  a  solu- 
tion of  ammonium  molybdate,  (NH4)2MoO4,  which  con- 
tains nitric  acid,  it  gives  a  complicated  insoluble  salt, 
ammonium  phospho-molybdate  (which  see). 

Pyrophosphoric  Acid,  H4P2O7.  —  When  phosphoric  acid 
is  heated  to  200°-300°  until  a  specimen  neutralized  with 
ammonia  gives  a  pure  white  precipitate  with  silver  nitrate. 
it  is  completely  transformed  into  pyrophosphoric  acid  by 


METAPHOSPHORIC  ACID.  331 

loss  of  water.  The  white  precipitate  referred  to  is  the 
silver  salt  of  pyrophosphoric  acid.  The  silver  salt  of 
orthophosphoric  acid  is  yellow.  This  difference  in  color 
led,  many  years  ago,  to  a  careful  investigation  of  the 
change  in  composition  which  phosphoric  acid  undergoes 
when  heated,  and  to  the  recognition  of  the  existence  of 
pyrophosphoric  acid  as  distinct  from  orthophosphoric 
acid ;  and  the  study  of  the  relations  existing  between 
these  acids  and  metaphosphoric  acid  has  had  a  strong 
influence  in  shaping  the  views  of  chemists  in  regard  to 
the  relations  between  other  similar  acids.  The  views  at 
present  held  in  regard  to  the  relations  between  the  com- 
mon forms  of  oxygen  acids  and  the  so-called  normal  acids 
or  maximum  hydroxides  are  simply  an  extension  of  the 
ideas  first  introduced  into  chemistry  in  connection  with 
the  three  varieties  of  phosphoric  acid.  The  different 
varieties  of  periodic  acid,  and  the  modifications  of  sul- 
phuric acid  seen  in  the  normal  acid,  S(OH)6,  the  ordinary 
acid,  SO2(OH)2,  and  the  pyro-acid,  H2S2O7,  are  examples 
of  the  same  kind  of  relations. 

Pyrophosphates  are  formed,  as  we  have  seen,  when 
the  secondary  phosphates,  like  disodium  phosphate, 
HNa2PO4,  are  heated. 

Metaphosphoric  Acid,  HPO3. — This  acid  is  formed  by 
dissolving  phosphorus  pentoxide,  P2O5,  in  cold  water : 

P205  +  H20  =  2HP03. 

It  is  also  formed  by  heating  phosphoric  acid  to  400°  : 
H3P04  =  HP03  +  H2O. 

Further,  the  metaphosphates  are  formed  by  heating 
the  primary  phosphates  like  primary  sodium  phosphate, 
H.2NaPO4.  The  acid  is  a  vitreous  translucent  mass,  known 
in  the  market  as  glacial  phosphoric  acid  (Acidum  phos- 
pJioricum  glaciale).  It  is  the  more  common  commercial 
form  of  phosphoric  acid.  It  is  a  monobasic  acid,  and  in 
composition  is  analogous  to  nitric  and  chloric  acids  : 

HPO3,     .     .     .     c     .     Metaphosphoric  acid. 

HNO3, Nitric  acid. 

HC1O3,    .     .     .     .     .     Chloric  acid. 


332  INORGANIC  CHEMISTRY. 

When  boiled  with  water  in  which  there  is  a  little  nitric 
acid  metaphosphoric  acid  is  readily  converted  into  ortho- 
phosphoric  acid : 

HPO3  +  H20  =  H3P04. 

This  transformation  is  effected  also  by  simply  allowing 
the  solution  of  the  meta-acid  in  water  to  stand  for  a  time, 
and  by  boiling  the  solution. 

When  a  metaphosphate,  as,  for  example,  sodium  meta- 
phosphate,  NaPO3,  is  heated  in  contact  with  a  metallic 
oxide,  it  takes  up  the  oxide  as  the  free  acid  takes  up 
water,  and  phosphates  are  thus  formed  in  which  two  or 
more  metals  take  the  place  of  the  hydrogen  of  the  acid. 
With  a  metallic  oxide  of  the  formula  M2O  it  combines 
to  form  a  phosphate,  M2NaPO4,  thus : 

NaPO3  +  M20  =  M2NaPO4 

— a  kind  of  action  which  is  plainly  analogous  to  the  con- 
version of  metaphosphoric  into  orthophosphoric  acid. 
So,  also,  when  an  oxide  of  the  formula  MO  is  heated 
with  sodium  metaphosphate  a  phosphate  of  the  formula 
MNaPO4,  in  which  M  represents  a  bivalent  metal,  is 
formed : 

NaP03  +  MO  =  MNaPO4. 

Upon  facts  of  this  kind  depends  the  power  of  sodium 
metaphosphate  to  dissolve  metallic  oxides,  as  when  beads 
formed  by  heating  sodium  ammonium  phosphate  are 
used  in  analysis.  The  first  effect  of  heating  the  phos- 
phate is,  as  explained  above,  the  formation  of  sodium 
metaphosphate  which  melts,  forming  a  clear  liquid  known 
as  the  "  bead  of  microcosmic  salt." 

Phosphorous  Acid,  H3PO3. — This  acid  is  formed  when 
pliosphorus  trichloride  is  treated  with  water.  It  is  also 
formed  together  with  phosphoric  and  hypophosphoric 
.acids  when  phosphorus  is  allowed  to  lie  in  contact  with 
moist  air.  The  acid  can  be  obtained  from  its  solutions 
by  evaporation,  when  it  is  deposited  in  transparent  crys- 
tals. When  heated  it  is  converted  into  phosphoric  acid, 
phosphine  being  given  off : 

4H3P03  =  3H3PO4  +  PH3. 


HYPOPHOSPHORIC  AND  HYPOPHOSPHOROUS  ACID.     333 

This  reaction  has  been  discussed  under  Phosphine  (which 
see).  The  tendency  of  phosphorous  acid  to  take  up 
oxygen  and  form  phosphoric  acid  makes  it  a  good  reduc- 
ing agent.  Its  action  is  well  illustrated  in  the  case  of 
mercuric  chloride,  HgCl2,  which  it  transforms  into  mer- 
curous  chloride,  HgCl.  Water  being  present,  the  phos- 
phorous acid  appropriates  the  oxygen  of  a  part  of  it, 
leaving  the  hydrogen  to  act  upon  the  mercuric  chloride  : 

2HgCl2  +  H3P03  +  H20  =  H3P04  +  2HgCl  +  2HC1. 

Phosphorous  acid  is  only  dibasic,  its  salts  having  the 
general  formula  HM2PO3.  This  fact  has  led  to  the  be- 
lief that  in  the  acid  two  of  the  hydrogen  atoms  are  in 
combination  with  oxygen  in  the  form  of  hydroxyl,  while 
the  third  is  in  combination  with  phosphorus  as  repre- 

(H 
sented  in  the  formula  OP  <  OH .     This  conclusion  finds 

(OH 

further  support  in  the  conduct  of  some  derivatives  of 
phosphorous  acid. 

Hypophosphoric  Acid,  H4P2O6,  is  formed  together  with 
phosphoric  and  phosphorous  acids  when  sticks  of  ordi- 
nary phosphorus  placed  in  glass  tubes  drawn  out  to  a  small 
opening  at  one  end  are  exposed  to  the  action  of  moist  air. 
By  arranging  a  number  of  such  tubes  on  a  funnel  the 
lower  end  of  which  is  in  a  bottle,  a  solution  is  gradually 
collected  which  contains  hypophosphoric  acid  together 
with  the  other  two  acids  mentioned.  The  salts  o£  the 
acid  show  that  it  is  tetrabasic.  It  has  been  suggested 
that  this  acid  has  the  constitution  represented  by  the 

OP(OH), 
formula       i 

OP(OH), 

Hypophosphorous  Acid,  H3PO2,  has  already  been  re- 
ferred to,  as  its  potassium  salt  is  formed  in  the  prepara- 
tion of  phosphine  by  the  action  of  phosphorus  upon  a 
solution  of  potassium  hydroxide  : 

3KOH  +  4P  +  3H20  =  3KH2PO2  +  PH3. 

The  acid  is  a  solid  which  crystallizes  well.  The  most 
characteristic  fact  in  its  conduct  is  its  marked  tendency 


334  INORGANIC  CHEMISTRY. 

to  pass  over  into  phosphoric  acid  by  taking  up  oxygen. 
It  is  therefore  a  good  reducing  agent.  It  reduces  sul- 
phuric acid  to  sulphurous  acid  and  even  to  sulphur,  as 
represented  in  the  two  equations  below : 

2H2SO4  +  H3P02  =  2H2S03     +  H3PO4 ; 
H,S03  +  H3P02  =  S  +  H20  +  H3PO, 

When  heated,  also,  it  forms  phosphoric  acid  and  phos- 
phine  just  as  phosphorous  acid  does  : 

mfO,  =,  HaPO,  +  PH3. 

The  acid  is  monobasic,  and  this  has  led  to  the  belief  that 
only  one  of  the  hydrogen  atoms  in  the  molecule  of  the 
acid  is  in  combination  with  oxygen  as  hydroxyl,  and  that 
the  two  others  are  in  combination  with  phosphorus  as 

(H 

represented  in  the  formula  OP-<  H   .     The  relation  be- 

(OH 

tween  this  acid  and  phosphorous  and  phosphoric  acids 
has  already  been  commented  upon  (see  page  323). 

Phosphorus  Pentoxide,  Phosphoric  Anhydride,  P2O5. — 
This  highest  oxidation-product  of  phosphorus  is  formed 
by  burning  the  element  in  air  or  in  oxygen.  It  is  a 
white  powder  which  attracts  moisture  from  the  air  and 
becomes  liquid.  This  power  to  combine  with  water  is 
its  most  characteristic  property.  It  forms  first,  as  we 
have  seen,  metaphosphoric  acid  and,  by  further  action, 
orthophosphoric  acid.  Its  action  towards  water  is 
strongly  suggestive  of  the  action  of  sulphur  trioxide 
or  sulphuric  anhydride  towards  water.  Owing  to  this 
power  to  combine  with  water,  phosphorus  pentoxide  is 
used  for  the  purpose  of  drying  gases,  and  as-  a  dehy- 
drating agent. 

Phosphorus  Trioxide,  or  Phosphorous  Anhydride,  P2O3 
(or  P4O6),  is  formed  by  burning  phosphorus  so  that  the 
air  does  not  have  free  access  to  it,  as  by  putting  a  piece 
of  phosphorus  in  a  glass  tube  drawn  out  to  a  fine  open- 
ing, drawing  air  over  the  phosphorus,  and  warming  it 
gently.  In  this  way  not  enough  air  can  get  access  to  the 


CONSTITUTION  OF  TEE  ACIDS  OF  PHOSPHORUS.  335 

phosphorus  to  convert  it  into  the  pentoxide.  The  tri- 
oxide  has  such  a  strong  tendency  to  pass  over  into  the 
pentoxide  that  when  brought  into  the  air  it  takes  fire  and 
burns,  forming  the  higher  oxide.  It  is  readily  converted 
into  phosphorous  acid  by  water. 

Phosphorus  Suboxide,  P^O,  is  one  of  the  products 
formed  by  the  burning  of  phosphorus  in  a  limited  sup- 
ply of  air.  It  is  also  formed  in  several  other  ways,  and 
two  varieties  of  it  have  been  described. 

Phosphorus  Tetroxide,  P2O4,  is  also  formed  when  phos- 
phorus is  burned  in  the  air.  It  can  be  obtained  pure  in 
the  form  of  colorless  crystals. 

Constitution  of  the  Acids  of  Phosphorus. — Considerable 
has  already  been  said  on  this  subject  in  dealing  with  the 
relations  between  the  acids.  The  view  that  phosphoric 
acid  contains  three  hydroxyl  groups  is  based  upon  the 
fact  that  the  acid  is  tribasic,  which,  taken  together  with 
what  is  known  in  regard  to  the  conduct  of  other  acids, 
suggests  that  all  three  hydrogen  atoms  in  the  molecule 
are  in  combination  with  oxygen.  This  view  is  the  sim- 
plest, and  all  facts  known  in  regard  to  the  conduct  of 
phosphoric  acid  are  in  accordance  with  it.  The  constitu- 

/0-H 
tion  is  represented  by  the  formula  O=P— O-H ,  which 

\O-H 

may  also  be  written  in  this  way  :  OP(OH)3.  Two  views 
suggest  themselves  in  considering  the  constitution  of 
phosphorous  acid.  It  may  be,  like  phosphoric  acid,  a  tri- 
hydroxyl  derivative  of  the  formula  P(OH)3,  or  it  may  have 

/H 

the  structure  represented  by  the  formula  O=P^-OH    or 

\OH 

(  H 

OP  -j  /QTT\  •  The  easy  formation  of  the  acid  from  phos- 
phorus trichloride  and  water  is  in  accordance  with  the 
former  view.  On  the  other  hand,  as  has  already  been 
remarked,  the  fact  that  the  acid  is  dibasic  speaks  against 
this  view,  and  in  favor  of  the  latter.  A  somewhat  com- 
plex reaction  of  an  organic  derivative  of  phosphorous  acid 
also  furnishes  evidence  in  favor  of  the  view 'that  there  are 
only  two  hydroxyl  groups  contained  in  the  molecule  of 


336  INORGANIC  CHEMISTRY. 

phosphorous  acid,  and  that  its  structure  is  represented  by 

/H 
the  formula  O-P-O-H. 


Phosphorus  Oxy  chloride,  POC13.  —  This  compound  has 
been  referred  to  in  connection  with  the  chlorides  of  phos- 
phorus. It  is  formed  by  the  action  of  ozone  on  phos- 
phorus trichloride  and  by  the  action  of  water  upon  the 
pentachloride  : 

PC13  +  0       =  POC13  ; 

PC1B  +  H3O  =  POC13  +  2HC1. 

It  may  be  regarded  as  phosphoric  acid  in  which  all  three 
of  the  hydroxyl  groups  are  replaced  by  chlorine,  just  as 
sulphuryl  chloride,  SO2C12,  is  to  be  regarded  as  sulphuric 
acid  in  which  both  hydroxyls  are  replaced  by  chlorine". 
The  fact  that  when  treated  with  water  and  other  com- 
pounds containing  hydroxyl  it  yields  phosphoric  acid  has 
been  mentioned,  and  the  value  of  this  reaction  and  the 
similar  reaction  of  phosphorus  pentachloride  as  a  means 
of  detecting  the  hydroxyl  condition  in  compounds  has 
been  pointed  out  (see  p.  316). 

Arsenic  Acid,  H3AsO4.—  The  compound  of  arsenic  and 
oxygen  which  is  most  readily  obtained  is  the  trioxide, 
As2O3,  and  this  is  formed  by  direct  combination  of  the  two 
elements.  When  this  is  oxidized  either  with  aqua  regia 
or  by  passing  chlorine  into  water  in  which  the  trioxide  is 
suspended  it  is  converted  into  arsenic  acid  : 

As2O3  +  3H20  +  2O  =  2H3AsO4. 

From  its  solutions  it  is  obtained  in  crystallized  form. 
According  to  the  temperature  to  which  it  is  heated 
the  deposit  has  the  composition  of  the  ortho-acid,  H3AsO4, 
of  the  pyro-acid,  H4As2O7,  or  of  the  meta-acid,  HAsO3. 
Perfect  analogy  with  the  phosphorus  compounds  is  here 
observed.  When  the  pyro-  and  meta-acids  are  dis- 
solved in  water  they  pass  at  once  into  the  form  of  the 
ortho-acid.  Arsenic  acid,  like  phosphoric  acid,  is  a  strong 
tribasic  acid,  forming  three  series  of  salts  which  under 


ARSENIOUS  ACID.  337 

the  influence  of  heat  conduct  themselves  like  the  cor- 
responding phosphates,  the  primary  salts  yielding  pyro- 
arsenates,  and  the  secondary  salts  yielding  meta-arsen- 
ates.  When  these  are  dissolved  in  water  they  pass  at 
once  into  the  corresponding  salts  of  ortho-arsenic  acid. 
Arsenic  acid  is  easily  reduced  to  the  form  of  arsenic. 
When  hydrogen  sulphide  is  passed  through  a  hydro- 
chloric acid  solution  of  arsenic  acid  different  reactions 
take  place  according  to  the  conditions.  The  three  pos- 
sibilities are  :  (1)  The  formation  of  the  pentasulphide  ; 

(2)  the  formation  of  sulphoxyarsenic  acid,  H3AsO3S  ;  and 

(3)  the  formation  of  arsenic  pentasulphide,  arsenic  trisul- 
phide,  and  sulphur.     These  reactions  are  represented  by 
the  following  equations  : 

(1)  H3As04      +  H2S     =  H3AsO3S  +  H2O  ; 

(2)  2H3As03S  +  3H2S  =  As2S5        +  6H2O  ; 


,Qx    j  2H3As03S  +  6HC1  =  2AsCl3      +  6H2O  +  28  ; 

I 


33  3  2 

2AsCl3       +  3H2S  =  As2S3        +  6HC1. 


The  first  action  is  that  represented  by  equation  (1). 
The  acid  thus  formed,  known  as  sulphoxyarsenic  acid, 
differs  from  arsenic  acid  only  in  the  fact  that  it  contains 
a  sulphur  atom  in  the  place  of  one  oxygen  atom.  It  is 
soluble  in  water,  and,  therefore,  when  hydrogen  sulphide 
is  passed  into  a  solution  of  arsenic  acid  there  is  at 
first  no  precipitate  formed  ;  but  gradually,  where  the  hy- 
drogen sulphide  is  in  excess,  some  of  the  sulphoxyar- 
senic acid  is  changed  to  arsenic  pentasulphide,  while  an- 
other part  of  the  acid  is  decomposed  by  hydrochloric 
acid,  forming  arsenic  chloride  and  sulphur,  and  the  tri- 
sulphide  is  then  precipitated.  Therefore,  the  precipitate 
formed  by  passing  hydrogen  sulphide  into  a  solution 
of  arsenic  acid  is  likely  to  consist  of  a  mixture  of  arsenic 
pentasulphide,  trisulphide,  and  sulphur. 

Arsenious  Acid,  H3AsO3,  is  not  known  in  the  free  state, 
but  salts  related  to  it  are  formed  by  treating  arsenic  tri- 
oxide  with  bases.  Thus,  when  it  is  treated  with  potas- 
sium hydroxide  the  salt  KAsO2,  or  potassium  meta- 
arsenite,  is  formed  : 

AsaO,  +  2KOH  =  2KAs02  +  H9O. 


338  INORGANIC  CHEMISTRY. 

Salts  of  meta-arsenious  acid,  AsO.OH,  are  more  com- 
monly obtained  than  those  of  the  normal  acid,  As(OH)3. 
In  alkaline  solution  arsenious  acid  tends  to  pass  into 
the  form  of  arsenic  acid,  and  it  is  therefore  a  useful 
reducing  agent.  Its  action  in  this  way  is,  however,  not 
as  strong  as  that  of  phosphorous  acid. 

Arsenic  Trioxide,  As2O3. — This  compound  is  commonly 
called  arsenic  or  white  arsenic.  It  is  the  most  important 
of  all  the  compounds  of  the  element  arsenic.  It  finds 
applications  for  many  purposes,  and  is  manufactured  in 
large  quantities.  It  occurs  in  small  quantity  in  nature, 
but  that  which  comes  into  the  market  is  manufactured 
by  roa-sting  natural  arsenides,  particularly  arsenical  py- 
rites, FeAsS.  The  products  of  roasting  this  compound 
are  ferric  oxide,  FeQO3,  sulphur  dioxide,  SO2,  and  arsenic 
trioxide,  As2O3.  Of  these,  the  first  is  a  non-volatile  solid, 
the  second  a  gas,  and  the  third  a  volatile  solid.  By  pass- 
ing the  volatile  products  through  properly  constructed 
canals  the  arsenic  trioxide  is  condensed  on  the  walls. 
Some  of  the  powder  thus  obtained  must  be  subjected  to 
a  second  process  of  distillation  to  make  it  pure  enough 
for  the  market.  In  a  recent  year  over  6000  tons  of  this 
substance  were  produced  in  England  and  Saxony. 

Arsenic  trioxide  is  a  colorless,  amorphous,  vitreous 
mass.  Gradually  it  becomes  opaque  and  crystalline,  with 
an  appearance  like  that  of  porcelain.  It  crystallizes  in 
two  forms,  the  common  one  being  that  of  regular  octa- 
hedrons. Under  exceptional  conditions  it  crystallizes  in 
the  form  of  rhombic  prisms.  When  heated  it  sublimes, 
and  is  deposited  on  a  cold  surface  in  the  form  of  octa- 
hedrons. Arsenic  trioxide  is  difficultly  soluble  in  water, 
but  more  easily  in  hydrochloric  acid.  The  solution  in 
hydrochloric  acid  contains  arsenic  trichloride  (see  p.  317), 
and  when  the  solution  is  boiled  the  chloride  is  carried 
over.  When  the  solution  of  the  amorphous  oxide  in 
hydrochloric  acid  is  concentrated  enough  it  deposits  the 
oxide  in  crystalline  form,  and  the  formation  of  the 
crystals  is  accompanied  by  an  evolution  of  light  which 
can  be  seen  in  a  dark  room.  When  the  crystalline 
variety  is  dissolved  it  is  deposited  in  crystals  without 


ARSENIC  TRIOXIDE.  339 

evolution  of  light.  The  formation  of  arsenic  trichloride 
by  the  action  of  hydrochloric  acid  on  the  oxide  is  perfectly 
analogous  to  the  formation  of  the  chloride  of  any  base- 
forming  element  by  the  action  of  hydrochloric  acid  upon 
the  oxide,  as,  for  example,  ferric  oxide.  The  reactions 
are  represented  thus : 

As2O3  +  6HC1  =  2AsCl3  +  3H2O ; 
and  Fe203  +  6HC1  =  2FeCl3  +  3H2O. 

While  in  this  reaction  arsenic  appears  as  a  base-forming 
element,  its  character  as  an  acid-forming  element  shows 
itself  when  the  chloride  is  treated  with  a  large  excess  of 
water,  under  which  circumstances  it  is  completely  con- 
verted into  the  oxide.  Towards  some  acids  also  arsenic 
trioxide  acts  as  a  weak  base.  A  somewhat  complex  sul- 
phate is  known  in  which  the  arsenic  replaces  a  part  of 
the  hydrogen  of  the  acid.  It  is  formed  by  treating  the 
trioxide  with  fuming  sulphuric  acid. 

The  trioxide  is  easily  reduced.  When  heated  with 
potassium  cyanide,  KCN,  or  with  charcoal  in  a  dry  glass 
tube  arsenic  is  deposited  above  the  flame  in  the  form  of 
a  dark  lustrous  layer.  When  brought  into  a  vessel  from 
which  hydrogen  is  being  evolved  it  is  reduced  to  arsine. 

The  specific  gravity  of  the  vapor  of  the  oxide  shows 
that  it  has  the  formula  As4O6,  and  not  As2O3 ;  as,  however, 
most  of  its  reactions  can  be  more  conveniently  expressed 
by  the  aid  of  the  simpler  formula,  the  latter  is  commonly 
used. 

Arsenic  trioxide  has  a  weak,  disagreeable,  sweet  taste, 
and  is  an  active  poison.  A  dose  of  from  two  to  three 
grains  is  sufficient  to  cause  death  unless  it  is  ejected  by 
vomiting,  or  rendered  harmless  by  being  converted  into 
an  insoluble  compound.  It  is  possible,  by  beginning  with 
small  doses,  and  gradually  increasing  them,  to  accus- 
tom the  human  body  to  considerably  larger  doses 
than  that  mentioned.  It  strengthens  the  power  of  the 
respiratory  organs,  and  consequently  facilitates  mountain- 
climbing.  The  peasants  in  some  mountain  regions  are 
said  to  use  it  habitually.  It  is  much  used  in  medicine, 
especially  in  skin  diseases.  It  is  also  used  extensively 


340  INORGANIC  CHEMISTRY. 

as  a  rat-poison.  The  most  efficient  antidote  is  a  mixture 
of  ferric  hydroxide,  Fe(OH)3,  and  magnesia,  which  forms 
with  arsenic  trioxide  an  insoluble  compound. 

Arsenic  Pentoxide,  As2O5,  is  formed  by  igniting  arsenic 
acid.  If  heated  too  high  the  pentoxide  breaks  down  into 
arsenic  trioxide  and  oxygen.  A  marked  difference  will  be 
observed  between  the  conduct  of  the  oxides  of  phosphorus 
and  that  of  the  corresponding  oxides  of  arsenic.  While 
phosphorus  trioxide  takes  up  oxygen  spontaneously  when 
exposed  to  the  air,  and  the  pentoxide  is  not  decomposed 
by  heat,  the  trioxide  of  arsenic  does  not  under  any  cir- 
cumstances take  up  oxygen  directly,  and  the  pentoxide 
easily  breaks  down  into  the  trioxide  and  oxygen  when 
heated. 

Sulphides. — There  are  three  compounds  of  arsenic  with 
sulphur — the  disulphide,  As2S2,  the  trisulphide,  As2S3, 
and  the  pentasulphide,  As2S5. 

Arsenic  Disulphide,  As2S2,  occurs  in  nature  and  is 
known  as  realgar.  It  can  also  be  obtained  by  melting 
arsenic  and  sulphur  together  in  the  right  proportions. 
It  forms  an  orange-red  powder  which  was  formerly  used 
as  a  pigment. 

Arsenic  Trisulphide,  As2S3,  is  found  in  nature  and  is 
called  orpiment  or  king's  yellow.  It  can  be  prepared  by 
melting  together  arsenic  and  sulphur  in  the  proper  pro- 
portions, and  by  precipitating  a  solution  of  arsenic  tri- 
oxide in  hydrochloric  acid  with  hydrogen  sulphide.  It 
melts,  forming  a  red  liquid.  The  natural  substance,  as 
well  as  that  which  is  precipitated  by  means  of  hydrogen 
sulphide,  is  yellow.  It  dissolves  in  soluble  sulphides, 
forming  salts  of  sulpharsenious  acid,  H3AsS3,  or  HAsS2. 
The  salts  are,  for  the  most  part,  derived  from  the  acid  of 
the  latter  formula.  There  is,  therefore,  perfect  analogy 
between  the  oxygen  and  sulphur  compounds,  for,  as  we 
have  seen,  when  arsenic  trioxide  is  dissolved  in  potassium 
hydroxide  a  salt  of  the  formula  KAsO2  is  formed.  The 
analogy  is  clearly  shown  by  means  of  the  equations 

As2O8  +  2KOH  =  2KAsO2  +  H2O  ; 

AsaS3  +  2KSH  =  2KAsS2  +  H2S. 


ARSENIC  TRISULPHIDE.  341 

The  acid  HAsS2  is  derived  from  the  corresponding  nor- 

(SH 
mal  acid  As  -<  SH  ,  by  loss  of  one  molecule  of  hydrogen 

(  SH 
sulphide : 

(SH 

As^  SH  = 

(SH 

just  as  the  acid  HAsO2  is  derived  from  the  normal  oxy- 

(OH 
gen  acid  As  -<  OH ,  by  loss  of  one  molecule  of  water  : 

(OH 

(OH 

As^OH  = 
(OH 

When  a  solution  of  a  sulpharsenite  is  treated  with  one 
of  the  stronger  acids,  as,  for  example,  hydrochloric  acid, 
arsenic  trisulphide  is  precipitated.  We  should  naturally 
look  for  the  separation  of  the  free  acid  according  to  the 
equation 

KAsS2  +  HC1  =  HAsS2  +  KC1 ; 

but,  if  this  is  formed,  it  breaks  down  at  once  into  hydro- 
gen sulphide  and  arsenic  trisulphide  : 

2HAsS2  =  As2S3  +  H2S. 

There  is  a  striking  analogy  between  this  action  and  that 
which  takes  place  when  a  stronger  acid  is  added  to  a 
solution  of  a  carbonate,  when  carbon  dioxide  is  set  free : 

K,C03  +  2HC1  =  H,C03  +  2KC1 ; 
H2C03  =  C02  +  H20. 

A  marked  difference  between  the  two  cases  is  to  be 
found  in  the  fact  that  the  trisulphide  of  arsenic  is  insol- 
uble in  water  and  therefore  appears  as  a  precipitate, 
while  carbon  dioxide  escapes  as  a  gas. 

Besides  salts  of  the  acids  H3AsSs  and  HAsS2 ,  there  are 
others  derived  from  the  more  complex  acid  H4As3S6. 


342  INORGANIC  CHEMISTRY. 

This  bears  to  normal  sulpharsenious  acid,  As(SH)3,  a  re- 
lation similar  to  that  which  pyrophosphoric  acid  bears 
to  orthophosphoric.  If  two  molecules  of  the  normal  acid 
lose  one  molecule  of  hydrogen  sulphide,  this  pyrosulph- 
arsenious  acid  is  the  product : 

(SH 

2As-{  SH  =  As2S(SH)4  +  H2S. 
(SH 

It  is  a  salt  of  this  acid  which  is  formed  when  arsenic 
trisulphide  is  dissolved  in  ammonium  sulphide : 

As2S3  +  2(NH4)2S  =  As2S(SNH4)4. 

Arsenic  Pentasulphide,  As2S3,  is  formed  by  melting  sul- 
phur and  arsenic  together  in  the  proper  proportions, 
and  by  precipitating  a  solution  of  sodium  sulpharsenate 
with  hydrochloric  acid  : 

2Na3AsS4  +  6HC1  =  6NaCl  +  As2S5  +  3H2S. 

Sulpharsenic  acids  corresponding  to  the  oxygen  acids 
suggest  themselves.  We  might,  for  example,  expect  to 
find  salts  derived  from  the  acids  H3AsS4,  HAsS3,  and 
H4As2S7,  corresponding  to  ortho-,  meta-,  and  pyro-arsenic 
acids.  When  arsenic  pentasulphide  is  dissolved  in  solu- 
tions of  metallic  sulphides  the  products  are  generally 
salts  of  pyrosulpharsenic  acid,  H4As2S7,  and  these  under- 
go decomposition  into  salts  of  the  ortho-  and  meta-acids. 
When,  for  example,  arsenic  pentasulphide  is  dissolved  in 
ammonium  sulphide  reaction  takes  place  thus : 

AsA  +  2(NH4)ZS  =  (NH4)4As,S, 

The  ammonium  salt  formed  in  this  way  is,  however,  de- 
composed thus : 

(NH4)4As,S,  =  (NH4),AsS4  +  (NH4)AsSs. 

Only  one  compound  intermediate  between  arsenic  and 
sulpharsenic  acids  is  known.  This  is  the  sulphoxyarsenic 
acid  formed  as  the  first  product  of  the  action  of  hydro- 


ANTIMONIC  ACID— ANTIMONY  TRIOXIDE.          343 

gen  sulphide  upon  a  solution  of  arsenic  acid,  which  was 
referred  to  under  Arsenic  Acid  (p.  337).  The  possibility 
of  other  products  of  the  formulas  H3AsO2S2  and 
H3AsOS3  will  occur  to  every  one. 

Antimonic  Acid,  H3SbO4. — This  acid  is  the  final  product 
of  the  oxidation  of  antimony  when  treated  with  aqua 
regia.  It  need  only  be  said  that  it  is  very  similar  to 
phosphoric  and  arsenic  acids ;  and  that,  like  these,  it 
yields  a  meta-  and  a  pyro-acid  of  the  formulas  HSbO3 
and  H4Sb2O7.  The  acid  of  the  formula  OSb(OH)3,  or 
orthoantimonic  acid,  is  known  in  the  free  state,  and  is 
formed  by  treating  a  soluble  salt  of  antimonic  acid  with 
sulphuric  or  nitric  acid  : 

OSb(OK)3  +  3HNO3  =  3KNO3  +  OSb(OH)3. 

An  acid  Sb2O(OH)8  is  also  known  in  the  free  state,  be- 
ing formed  by  the  action  of  antimony  pentachloride  upon 
water.  The  lower  oxides  of  antimony,  the  trioxide,  Sb2O3, 
and  the  tetroxide,  Sb2O4,  are  not  strongly  acidic;  that  is 
to  say,  they  do  not  readily  form  salts  when  treated  with 
bases.  In  this  respect  the  trioxide  of  antimony  differs 
markedly  from  the  corresponding  oxides  of  phosphorus 
and  arsenic. 

Antimony  Trioxide,  Sb2O3. — This  compound  is  found 
in  nature  as  white  ore  of  antimony,  and  is  easily  formed 
by  burning  antimony  in  the  air  and  by  oxidizing  it  with 
nitric  acid  or  saltpeter.  That  formed  by  burning  anti- 
mony in  the  air  always  contains  some  of  the  tetroxide,  and 
by  heating  it  long  enough  in  the  air  and  to  a  temperature 
high  enough  it  is  completely  transformed  into  the  tetrox- 
ide. When  the  trioxide  is  dissolved  in  caustic  soda  a 
salt  of  the  formula  NaSbO2  is  formed.  This  is  plainly 
derived  from  an  acid  of  the  formula  HSbO2,  which  bears 
a  simple  relation  to  normal  antimonious  acid.  Towards 
most  bases,  however,  antimony  trioxide  does  not  conduct 
itself  as  an  acid.  On  the  other  hand,  towards  the  stronger 
acids  it  acts  as  a  base. 

Salts  of  Antimony. — The  salts  of  antimony  are  derived 
either  from  the  hydroxide  Sb(OH)3,  or  from  the  hydrox- 
ide SbO.OH.  The  salts  of  the  first  class  are  called  anti- 


344  INORGANIC  CHEMISTRY. 

mony  salts  /  those  of  the  second  class  are  called  antimonyl 
salts.  In  the  salts  formed  when  the  trihydroxide  of  an- 
timony is  completely  neutralized  by  acids,  the  antimony 
takes  the  place  of  three  atoms  of  hydrogen.  Thus,  the 
nitrate  has  the  formula  Sb(NO3)3 ;  the  sulphate  has  the 
formula  Sb3(SO4)3 ;  etc.  Besides  these  normal  salts  there 
are,  however,  basic  salts.  Thus  there  are  two  basic 

(OH  (OH 

nitrates  possible  of  the  formulas  Sb  -<  OH  and  Sb  -<  NO3. 

I  NO,  (  NO, 

The  formation  of  antimonyl  salts  may  be  illustrated  by 
the  sulphate.  This  may  be  regarded  as  formed  by  the 
action  of  sulphuric  acid  upon  the  hydroxide  SbO.OH, 
which  is  analogous  in  composition  to  the  acid  of  arsenic 
of  the  formula  AsO.OH : 

2SbO.OH 

The  product  is  antimonyl  sulphate.  The  weak  basic 
character  of  the  hydroxides  of  antimony  is  shown  by  the 
fact  that  many  of  its  salts  are  decomposed  by  water. 
The  salt  of  antimony  which  is  most  commonly  met  with 
is  the  so-called  tartar  emetic,  which  appears  to  be  an  anti- 
monyl potassium  salt  of  tartaric  acid.  Tartaric  acid  is  a 

(  OTT 
dibasic  acid  of  the  formula  C4H4O4  -j  QTT  .     When  one  of 

its  acid  hydrogen  atoms  is  replaced  by  potassium,  and 
the  other  by  the  antimonyl  group  SbO,  the  salt  thus 

formed  is  tartar  emetic,  C4H4O4  \  QTT-  .  It  is  also  pos- 
sible that  this  salt  may  be  derived  from  the  trihydroxide 
Sb(OH)3  by  replacement  of  one  hydrogen  atom  by  potas- 
sium, and  neutralization  of  the  rest  of  the  compound  by 
the  dibasic  tartaric  acid.  It  seems  more  probable,  how- 
ever, that  when  tartaric  acid  acts  upon  the  compound 

Sb  j  x-r£  '2  it  first  appropriates  the  potassium  atom,  form- 
ing acid  potassium  tartrate,  and  that  the  antimony  triox- 
ide  being  basic  is  neutralized  by  the  acid  tartrate.  To 
decide  between  the  two  views  is  at  present  impos- 
sible. 


OXIDES  AND  SULPHIDES  OF  ANTIMONY.        345 

Antimony  trioxide  dissolves  in  hydrochloric  acid, 
forming  the  trichloride,  and  this,  as  has  been  stated,  is 
decomposed  by  water  yielding  oxychloricies. 

Antimony  Tetroxide,  Sb2O4. — This  compound  is  most 
easily  obtained  by  igniting  antimonic  acid,  H3SbO4.  Two 
reactions  are  of  course  involved  : 

2H3Sb04  =  Sb205  +  3H20  ; 
Sb206       =  Sb204  +  O. 

It  is  also  formed  by  igniting  the  trioxide  in  the  air.  At 
ordinary  temperatures  the  tetroxide  is  white,  but  it  be- 
comes yellow  when  heated.  Towards  strong  acids  this 
oxide  acts  like  a  weak  base.  A  potassium  salt  of  the 
formula  K2Sb2O5  is  known,  which  is  derived  from  the 
acid  H2Sb2O5,  and  this  in  turn  from  the  simpler  acid 
SbO(OH)2  by  loss  of  water.  The  oxide  itself  is  regarded 
by  some  as  an  antimonyl  salt  of  metantimonic  acid, 
SbO2.OH,  of  the  formula  SbO2.O.SbO. 

Antimony  Pentoxide,  Sb2O5. — The  tetroxide  of  anti- 
mony does  not  combine  with  oxygen  to  form  the  pentox- 
ide.  The  latter  can  be  obtained  only  by  gentle  ignition 
of  antimonic  acid,  care  being  taken  not  to  raise  the  tem- 
perature high  enough  to  decompose  the  pentoxide  into 
the  tetroxide  and  oxygen.  The  fact  that  the  pentoxide 
readily  yields  salts  of  antimonic  acid  when  treated  with 
basic  solutions  was  mentioned  under  Antimonic  Acid. 

Antimony  Trisulphide,  Sb2S3. — This  compound  occurs 
in  nature  in  considerable  quantity  and  is  the  chief  source 
of  antimony.  It  is  known  as  stibnite  and  antimony  blende. 
In  some  localities,  especially  in  Japan,  it  occurs  in  large 
€rystals  of  great  beauty.  When  heated  in  the  air,  or 
roasted,  it  is  converted  into  the  trioxide,  and  finally  into 
the  tetroxide,  while  the  sulphur  escapes  as  the  dioxide. 
Hydrochloric  acid  dissolves  the  trisulphide  in  the  form 
of  the  chloride  with  evolution  of  hydrogen  sulphide  : 

Sb2S3  +  6HC1  =  2SbCl3  +  3H2S. 

Nitric  acid  converts  it  into  the  oxide  with  separation  of 
sulphur.  When  a  solution  of  antimony  chloride  is 
treated  with  hydrogen  sulphide,  the  trisulphide  is  thrown 


346  INORGANIC  CHEMISTRY. 

down.  This  artificially  prepared  trisulphide  has  an 
orange-red  color,  while  that  which  occurs  in  nature  is 
black  or  gray.  The  sulphide  dissolves  in  solutions  of 
metallic  sulphides,  forming  salts  of  sulphantimonious 
acid,  either  SbS.SH  or  Sb(SH)3. 

Antimony  Pentasulphide,  Sb2S5,  is  formed  by  passing 
hydrogen  sulphide  into  a  solution  of  antiinonic  acid  or 
by  decomposing  a  salt  of  sulphantimonic  acid  by  means 
of  an  acid.  The  action  takes  place  thus  : 

2H3Sb04  +  5H2S  =  Sb2SB  +  8H20  ; 
2Na,SbS4  +  6HC1  =  GNaCl  +  Sb2S&  +  3H2S. 

It  is,  when  dry,  a  golden-yellow  powder  known  as  sul- 
phur auratum.  It  dissolves  easily  in  solutions  of  metallic 
sulphides,  forming  the  sulphantimonates,  of  which  the 
sodium  salt,  Na3SbS4,  known  as  Schlippe's  salt,  is  a  good 
example.  The  action  is  represented  by  this  equation  : 

Sb2S&  +  GNaSH  =  2Na3SbS4  +  3H2S. 

When  heated  in  the  air  the  pentasulphide  gives 
off  enough  sulphur  to  form  the  trisulphide ;  while 
when  the  pentoxide  is  heated  it  is  converted  into  the 
tetroxide.  The  sulphantimonates  are  decomposed  when 
treated  with  acids  and  the  pentasulphide  is  thrown  down. 

Constitution  of  the  Acids  of  Arsenic  and  Antimony. — 
There  is,  in  general,  marked  analogy  between  the  com- 
pounds of  phosphorus  and  those  of  arsenic  and  anti- 
mony. In  one  particular,  however,  there  is  a  difference 
which  is  worthy  of  special  mention.  It  appears  that, 
while  phosphorous  acid  is  dibasic  and  probably  has  the 

(H 
structure  OP  •<  OH ,  arsenious  and  antimonious  acids  are 

(OH 

the  normal  compounds  represented  by  the  formulas 
As(OH)3  and  Sb(OH)3.  Arsenic  and  antimonic  acids  ap- 
pear to  have  the  same  structure  as  phosphoric  acid 

represented  by  the  formulas  As  •]  /^TTX  and  Sb  -!  /ri-rjx  . 

(  (U±i)3  (  (O±i)3 

The  difference  between  phosphorous  and  arsenious  acids 
suggests  the  difference  between  sulphurous  and  selenious 


OXIDES  OF  BISMUTH.  347 

acids.   While,  according  to  the  evidence,  the  constitution 
of  sulphurous  acid  is  that  represented  by  the  formula 

{TT 
^TT  ,  that  of  selenious  acid  is  represented  by  the 

formulaOSe 


Oxychlorides  of  Antimony.  —  Under  the  head  of  Anti- 
mony Trichloride  the  fact  was  mentioned  that  this  com- 
pound is  decomposed  by  cold  water  as  represented  in 
the  equation 

SbCl3  +  H2O  =  SbOCl  +  2HC1. 

If,  however,  hot  water  is  used,  the  composition  of  the 
product  approximates  to  that  represented  by  the  formula 
Sb4O5Cl2.  This  complex  mixture  of  oxychlorides  is 
known  as  the  "Poivder  of  Algaroth"  It  may  be  regarded 
as  derived  from  the  simple  oxychloride  by  loss  of  anti- 
mony trichloride,  thus  : 

SSbOCl  =  Sb4OBCl2  +  SbCl3. 

Many  other  oxychlorides  besides  the  two  mentioned 
have  been  obtained,  but  they  are  all  more  or  less  closely 
related  to  the  simple  compound  SbOCl. 

Oxides  of  Bismuth.  —  The  principal  compound  of  bis- 
muth and  oxygen  is  the  trioxide,  Bi2O3,  which  is  formed 
when  bismuth  is  burned  in  the  air.  It  is  a  yellow  pow- 
der. Besides  the  method  just  mentioned,  it  is  formed 
by  decomposing  bismuth  nitrate  by  high  heat.  If  a  so- 
lution of  bismuth  nitrate,  Bi(NO3)3,  is  treated  with  a 
cold  solution  of  potassium  hydroxide,  bismuth  hydrox- 
ide, Bi(OH)3,  is  thrown  down.  When  this  is  dried  at 
100°  it  loses  water  and  is  converted  into  the  hydroxide, 
BiO(OH)  ;  and  if  the  hydroxide  first  precipitated  is  boiled 
with  the  solution  it  is  converted  into  the  yellow  oxide, 
Bi2O3.  The  reactions  involved  are  ,  . 

Bi(NO3)3  +  3KOH    =  Bi(OH)3  +  3KN03  ; 
Bi(OH)3  =  BiO.OH  +  H2O  ; 
2Bi(OH)3  =  Bi2O3       +  3H20. 

The  trioxide  of  bismuth  is  basic  and  forms  salts  which 
in  composition  correspond  to  the  salts  of  antimony. 


348  INORGANIC  CHEMISTRY. 

Like  the  latter,  they  are  of  two  classes  —  the  bismuth  salts 
and  the  bismuthyl  salts.  The  former  are  derived  from 
the  triacid  base,  Bi(OH)3,  the  latter  from  the  monacid 
base,  BiO(OH). 

Salts  of  Bismuth.  —  The  best  known  salts  of  bismuth 
are  those  which  it  forms  with  sulphuric  and  with  nitric 
acids.  There  is  a  sulphate  of  the  formula  BiH(SO4)3 
formed  by  dissolving  bismuth  oxide  in  dilute  sulphuric 
acid.  The  sulphate  which  is  most  stable  in  the  presence 
of  water  is  the  bismuthyl  salt,  (BiO)2SO4.  When 
bismuth  is  dissolved  in  nitric  acid  and  the  solution 
evaporated  to  dryness  the  salt  Bi(NO3)3  +  10H2O  is  ob- 
tained. This  salt  is  decomposed  when  heated,  and  by 
water,  forming  basic  nitrates  of  bismuth.  The  composition 
of  the  basic  nitrate  obtained  by  decomposing  the  neutral 
nitrate  with  water  differs  according  to  the  conditions. 
Hot  and  cold  water  produce  different  results.  A  solu- 
tion containing  much  nitric  acid  does  not  give  the  same 
result  as  one  which  contains  little,  etc.  As  basic  bismuth 
nitrate  is  used  in  medicine  it  is  necessary  that  specific 
directions  should  be  given  for  its  preparation,  in  order 
that  a  substance  of  the  same  composition  should  always 
be  obtained.  Among  the  basic  nitrates  which  have  been 


isolated   are   the    following  :    Bi  j  §^\  BiO.NO3  and 

(  O.BiO 
Bi-<  O.NO2  .    Besides  these  many  of  much  more  complex 

(OH 

composition  are  known,  but  all  of  them  can  be  referred 
to  the  simple  forms.  Some  of  them  are  of  special  inter- 
est, as  they  appear  to  be  derived  from  complex  forms  of 
nitric  acid,  as,  for  example,  an  acid  of  the  formula 
N2O3(OH)4  or  H4N2O7,  which  is  analogous  to  pyrophos- 
phoric,  pyroarsenic,  and  pyroantimonic  acids.  The  basic 
nitrate  of  bismuth,  or  the  subnitrate,  as  it  is  frequently 
called  in  pharmacy,  is  much  used  in  medicine  as  a  rem- 
edy in  dysentery  and  cholera.  It  is  also  used  as  a 
cosmetic. 

Bismuth  Dioxide,  Bi2O2,  is  formed  as  a  brown  precipi- 
tate when  potassium  hydroxide  is  added  to  a  solution  of 


COMPOUNDS  OF  BISMUTH.  349 

bismuth  chloride  and  stannous  chloride,  SnCl2.  Stan- 
nous  chloride  combines  very  readily  with  chlorine  to 
form  stannic  chloride,  SnCl4.  When,  therefore,  stannous 
chloride  and  bismuth  chloride  are  brought  together,  it 
is  probable  that  the  former  extracts  a  part  of  the  chlo- 
rine from  the  latter,  forming  a  chloride  of  the  formula 
BiCl2,  and  this  with  the  potassium  hydroxide  breaks 
down,  yielding  the  dioxide  : 

2BiCl3  +  Sn012    =  2BiCl2  +  SnCl4 ; 
2BiCla  +  4KOH  =  Bi202   +  4KC1  +  2H2O. 

Bismuth.  Pentoxide,  Bi2O5,  is  formed  by  oxidizing  the 
trioxide,  by  means  of  chlorine,  in  alkaline  solution.  Al- 
though some  experimenters  appear  to  have  obtained  salts 
of  bismuthic  acid,  as,  for  example,  KBiO3,  others  have 
failed  to  obtain  them.  In  any  case  it  is  evident  that  the 
acid  properties  of  the  oxide  are  very  weak. 

Bismuth.  Trisulphide,  Bi2S3,  occurs  in  nature,  and  is 
formed  by  precipitating  bismuth  from  solutions  of  its 
salts  with  hydrogen  sulphide.  It  dissolves  in  hot  con- 
centrated hydrochloric  acid  and  in  nitric  acid.  It  does 
not  dissolve  in  solutions  of  the  sulphides  as  the  sulphides 
of  arsenic  and  antimony  do. 

Bismuth  Oxychloride,  BiOCl,  which  in  composition  is 
analogous  to  the  simplest  form  of  antimony  oxy chloride, 
is  thrown  down  as  a  white  powder  when  a  solution  con- 
taining bismuth  chloride  is  treated  with  water : 

BiCl3  +  H3O  =  BiOCl  +  2HC1. 

FAMILY  Y,  GROUP  A. 

As  the  members  of  Group  A,  Family  VII,  are  related 
to  Group  B  of  the  same  family ;  and  as  the  members  of 
Group  A,  Family  VI,  are  related  to  the  members  of  Group 
B  of  the  same  family,  so  the  members  of  Group  A,.  Family 
V,  are  related  to  the  members  of  Group  B,  which  have  just 
been  studied.  The  members  of  Group  A  are  vanadium, 
columbium,  tantalum,  and  didymium,  all  of  which  are 
rare.  Of  these  vanadium  has  been  most  thoroughly  in- 
vestigated, and  columbium  next, 


350  INORGANIC  CHEMISTRY. 

Vanadium,  V  (At.  Wt.  50.99).  —This  element  occurs  in 
nature  in  the  form  of  vanadates  or  salts  of  vanadic  acid, 
H3VO4,  which  is  analogous  to  phosphoric  acid.  The 
methods  employed  in  separating  the  element  from  its 
compounds  depend  upon  the  composition  of  the  com- 
pound. In  the  separation  advantage  is  frequently  taken 
of  the  fact  that  the  ammonium  salt  of  vanadic  acid  is 
difficultly  soluble  in  a  solution  of  ammonium  chloride. 
When  this  ammonium  salt  is  ignited  it  is  converted  into 
the  pentoxide  V2O6.  With  chlorine,  vanadium  forms  the 
compounds  YC12,  YC13,  and  YC14 ;  with  oxygen,  the 
compounds  V2O,  Y2O2,  Y2O3,  Y2O4,  and  Y2O5.  In  its  re- 
lations to  oxygen  it  suggests  nitrogen.  The  oxide,  Y2O4, 
conducts  itself  something  like  the  tetroxide  of  antimony. 
Towards  strong  bases  it  acts  like  an  acid,  forming  salts 
of  the  general  formula  Y4O7(OM)2.  (See  Antimony 
Tetroxide.) 

Vanadic  Acid,  H3VO4,  is  the  most  important  and  best 
known  of  the  compounds  of  vanadium.  It  is  the  final 
product  of  the  oxidation  of  vanadium,  and  bears  to  this 
element  the  same  relation  that  phosphoric,  arsenic,  and 
antimonic  acids  bear  to  phosphorus,  arsenic,  and  anti- 
mony. The  vanadates  are  derived  from  ortho-,  meta-, 
and  pyro-vanadic  acids,  though  the  most  stable  ones  are 
the  metavanadates,  MYO3.  The  free  metavanadic  acid  is 
known.  It  is  a  beautiful  golden-yellow  compound,  which 
may  be  used  as  a  substitute  for  gold  bronze.  An  oxy- 
chloride  of  the  formula  YOC13,  corresponding  to  phos- 
phorus oxychloride,  is  made  by  direct  addition  of  chlorine 
to  vanadium  dioxide. 

Tantalum,  Ta  (At.  Wt.  181.45). — Tantalum  occurs  in  the 
minerals  columbite  and  tantalite,  accompanied  by  nio- 
bium. With  the  members  of  the  chlorine  group  it  forms 
the  compounds  TaF6,  TaCl6,  TaBr5,  and  TaI6.  Tantalum 
fluoride  combines  easily  with  the  fluorides  of  other  metals 
forming  the  fluotantalates.  These  may  be  regarded  as 
salts  of  fluotantalic  acid,  which  are  derived  from  the  oxy- 
gen acids  by  replacement  of  a  part  or  all  of  the  oxygen  by 
fluorine.  Thus,  the  salt  K2TaF7  is  easily  obtained  by 
treating  tantalum  fluoride  with  a  solution  of  potassium 


BORON.  351 

fluoride.  This  is  a  salt  of  the  acid  H2TaF7  or  H4Ta2F14, 
which  is  analogous  to  the  oxygen  acid  H4Ta2O7.  With 
oxygen  it  forms  Ta2O4  and  Ta2O5.  The  latter  forms  the 
tantalates  with  bases.  When  tantalum  pentachloride  is 
decomposed  with  water  it  forms  the  acid  H4TaaO,  or 
pyrotantalic  acid : 

2TaCl5  +  7HaO  =  Ta203(OH)4  +  10HC1. 

The  tantalates  are  derived  from  the  meta-acid  HTa03, 
and  from  the  hexa-acid  H8Ta6O19,  which  is  derived  from 
the  ortho-acid  as  represented  in  this  equation : 

6H3TaO4  =  H8Ta6O19  +  5H3O. 

Columbium,  Cb  (At.  Wt.  93.7). — This  element,  which  is 
sometimes  called  niobium,  occurs  in  the  mineral  colum- 
bite.  It  forms  two  chlorides,  CbCl3  and  CbCl6,  and  a 
bromide  and  fluoride  corresponding  to  the  latter  chlo- 
ride. The  fluoride  readily  forms  fluocolumbates,  similar 
to  the  fluotantalates.  The  niobates  are-  derived  from  a 
number  of  forms  of  the  acid  which  are,  however,  closely 
related  to  the  ortho-acid  H3CbO4. 

Didymium  consists  of  two  very  similar  elements,  neo- 
dymium  and  praseodymium.  In  some  of  their  compounds 
they  show  a  resemblance  to  the  members  of  this  group. 
They  form,  for  example,  an  oxide  of  the  formula  Di2O6. 
On  the  other  hand,  they  seem  to  be  more  closely  related 
to  cerium  and  lanthanum,  which  are  also  very  rare  ele- 
ments, occurring  associated  with  didymium.  These  will 
be  further  treated  of  in  connection  with  lanthanum  and 
cerium. 

BOEON,  B  (At.  Wt.  10.86). 

General. — Although  the  element  boron  is  not  a  mem- 
ber of  the  family  to  which  nitrogen  and  phosphorus  be- 
long, it  nevertheless  resembles  the  members  of  this 
family  in  some  respects.  It  belongs  to  the  same  family 
as  aluminium,  and  in  the  composition  of  its  compounds 
it  is  undoubtedly  similar  to  aluminium  ;  but,  on  the  other 
hand,  its  oxide  is  distinctly  acidic,  while  that  of  aluminium 
is  basic. 


352  INORGANIC  CHEMISTRY. 

Occurrence. — Boron  occurs  in  nature  chiefly  in  the  form 
of  boric  acid,  or  as  salts  of  this  acid,  particularly  a 
sodium  salt  known  as  borax. 

Preparation. — From  borax  and  the  other  borates  the 
acid  can  easily  be  obtained.  When  heated,  water  is  given 
off,  and  boron  trioxide,  B2O8 ,  is  left : 

2B(OH)3  =  B303  +  3H,0. 

By  heating  the  oxide  with  potassium  amorphous  boron 
is  obtained.  By  melting  the  oxide  with  aluminium, 
boron  is  formed  and  is  dissolved  in  the  molten  alumin- 
ium, from  which,  on  cooling,  it  is  deposited  in  crystals. 
Amorphous  boron  in  almost  pure  form  is  obtained  by 
heating  borax  with  magnesium  powder.  One  of  the 
chief  difficulties  encountered  in  preparing  boron  is  to 
prevent  the  element  from  combining  with  the  nitrogen 
of  the  air.  At  the  high  temperature  at  which  the 
reduction  takes  place  the  two  elements  combine  very 
readily  to  form,  the  compound  boron  nitride,  BN.  The 
crystals  obtained  in  the  process  described  are  not  pure 
boron,  but  contain  aluminium,  or  carbon  and  aluminium, 
apparently  in  combination  with  the  boron.  The  crystals 
are  very  hard,  and  some  of  them  have  a  high  lustre. 

Properties. — Amorphous  boron  is  a  greenish-brown 
powder.  It  burns  when  heated  in  the  air  or  in  oxygen, 
the  product  being  the  trioxide  B2O3.  Strong  oxidizing 
agents,  like  nitric  acid  and  saltpeter,  readily  oxidize  it> 
forming  boric  acid.  It  combines  readily  also  with  many 
other  elements,  as  with  chlorine,  nitrogen,  and  sulphur. 
When  it  is  brought  into  the  melting  hydroxides  or  car- 
bonates of  potassium  or  sodium,  it  forms  borates  of  the 
corresponding  metals. 

Boron  Trichloride,  BC19. — This  compound  is  formed  by 
heating  boron  in  a  current  of  dry  chlorine,  and  by  heat- 
ing a  mixture  of  boron  trioxide  and  charcoal  in  chlorine  : 

2B,08  +  30  +  601,  =  4BC13  +  3CO,. 

This  reaction  is  especially  interesting  on  account  of  its 
double  character.     Carbon  alone  could  not  reduce  the 


BORON  TRIFLUORIDE.  353 

boron  trioxide  at  the  temperature  employed  ;  nor  could 
the  chlorine  alone  displace  the  oxygen  and  form  the 
chloride,  but  when  both  chlorine  and  carbon  act  together 
these  changes  take  place,  one  aiding  the  other. 

The  chloride  is  a  liquid  which  boils  at  17°.  Like 
phosphorus  trichloride,  it  is  easily  decomposed  by  water, 
forming  boric  acid,  which,  as  will  be  seen,  is  analogous 
in  composition  to  phosphorous  acid  and  arsenious  acid : 

BC13  +  3H20  =  B(OH)3  +  3HC1. 

This  decomposition  is  analogous  to  that  of  arsenic  tri- 
chloride rather  than  to  that  of  phosphorus  trichloride, 
for  in  the  latter  case  a  secondary  change  takes  place,  re- 
sulting in  the  formation  of  an  acid  of  the  constitution 
H 


OP  -1  (OH); 

Boron  Trifluoride,  BF 3,  is  obtained  by  treating  a  mix- 
ture of  fluor-spar  and  boron  trioxide  with  concentrated 
sulphuric  acid.  The  reaction  is  a  double  one,  consisting, 
first,  in  the  setting  free  of  hydrofluoric  acid  from  the 
fluor-spar : 

CaF2  +  H2SO4  =  CaS04  +  2HF ; 

and,  second,  in  the  action  of  the  hydrofluoric  acid  upon 
the  oxide  of  boron : 

B2O3  +  6HF  =  2BF3  +  3H2O. 

It  is  a  colorless  gas,  which  acts  upon  water,  and  therefore 
forms  a  thick  white  cloud  in  the  air.  The  action  upon 
water  is  represented  by  the  equation 

4BF3  +  3H2O  =  B(OH)3  +  3HBF4. 

The  first  action  which  we  should  expect  is  the  formation 
of  normal  boric  acid,  thus  : 

BFS  +  3H2O  =  B(OH)3  +  3HF. 

But  the  hydrofluoric  acid  combines  with  some  of  the 
trifluoride  of  boron,  forming  the  compound  HBF4,  which 
is  known  as  fluoboric  acid.  Several  elements  act  in  this 


354  INORGANIC  CHEMISTRY. 

way,  particularly  the  members  of  the  silicon  group. 
Silicon  itself  forms  the  well-known  compound  fluosilicic 
acid.  Fluoboric  acid  is  to  be  regarded  as  metaboric  acid, 
HBO2,  in  which  the  two  oxygen  atoms  have  been  re- 
placed by  fluorine.  The  acid  has  been  obtained  in  the 
free  state,  and  is  a  liquid  boiling  at  120°.  It  forms  salts 
of  the  general  formula  MBF4,  of  which  the  potassium 
salt,  KBF4,  is  the  best  example. 

Boric  Acid,  B(OH)3. — Boric  acid  occurs  free  in  nature 
and  in  the  form  of  salts,  of  which  the  principal  one  is 
borax.  Besides  borax,  which  is  a  sodium  salt  derived 
from  tetraboric  acid,  H2B4O7,  there  are  other  natural 
borates,  as  boracite,  which  is  a  magnesium  salt  combined 
with  magnesium  chloride  ;  and  datholite,  which  is  made 
up  of  silicic  acid,  boric  acid,  and  the  element  calcium. 
One  of  the  most  interesting  natural  forms  of  boric  acid 
is  that  which  is  given  off  from  the  earth  with  steam. 
Such  jets  of  steam  are  met  with  in  many  volcanic  regions, 
and  are  called  fumaroles.  In  Tuscany  many  of  the 
fumarbles  are  charged  with  small  quantities  of  boric 
acid,  which  is  somewhat  volatile  with  steam.  Those  at 
Monte  Cerboli  and  Monte  Kotundo  in  Tuscany  are  util- 
ized for  the  purpose  of  obtaining  the  boric  acid.  For 
this  purpose  basins  are  built  over  the  fumaroles  and  filled 
with  water,  so  that  the  steam  is  condensed  and  the  boric 
acid  dissolved  in  the  water.  The  solutions  formed  at  the 
higher  levels  flow  into  basins  at  lower  levels,  and  finally 
become  charged  with  a  considerable  quantity  of  the  acid, 
when  it  is  evaporated  to  crystallization  by  the  aid  of  the 
heat  furnished  by  the  fumaroles.  The  acid  obtained  in 
this  way  is  not  pure,  but  by  recrystallization  it  is  easily 
purified. 

Boric  acid  can  also  be  made  from  borax  by  heating  the 
salt  in  solution  with  dilute  sulphuric  acid : 

Na3B40T  +  H2S04  +  5H3O  =  Na2SO4  +  4B(OH)3. 

If  the  solution  is  sufficiently  concentrated  the  boric  acid 
crystallizes  out  on  cooling. 

Boric  acid  is  easily  soluble  in  water,  and  crystallizes 
from  the  solution.  It  is  also  soluble  in  alcohol,  and  this 


BORIC  ACID.  355 

solution  burns  with  a  characteristic  green  flame.  The 
acid  is  quite  volatile  with  water  vapor.  When  heated  at 
100°  orthoboric  acid  loses  one  molecule  of  water,  and  is 
converted  into  metaboric  acid,  HBO2  ;  at  160°  it  yields 
tetraboric  acid,  H2B4O7  ;  and  at  a  higher  temperature  it  is 
converted  into  boron  trioxide  or  boric  anhydride,  BaO8. 
These  changes  are  represented  in  the  equations  follow- 
ing: 

OH  0 


/OH 

B(O 

M>H 


OH 


H 
OH 


\OH 

The  most  stable  salts  are  the  tetraborates  and  meta- 
borates.  Borax  is  the  sodium  salt  of  tetraboric  acid, 
Na2B4O7.  The  salts  of  orthoboric  acid  are  unstable. 
They  break  down  when  treated  with  water,  forming  free 
boric  acid  and  either  metaborates  or  tetraborates. 

When  heated  together  with  oxides,  boric  oxide  forms 
salts  just  as  boric  oxide  and  water  form  boric  acid. 
Borax  also,  when  treated  with  metallic  oxides,  forms 
double  borates,  which  are  derived  from  normal  boric 
acid.  Thus  with  copper  oxide  action  takes  place  which 
should  probably  be  represented  thus  : 

Na2B4O7  +  5CuO  =  Na2Cu6(BO3)4 ; 
and         BaO3        +  3CuO  =  Cu3(BO3)2. 

Many  of  these  salts  are  colored,  and  the  action  of 
metallic  compounds  upon  boron  trioxide  and  upon  borax 
ig  utilized  for  the  purpose  of  determining  their  nature 


356  INORGANIG  CHEMISTRY. 

by  the  color  of  the  mass  formed.  It  will  be  remembered 
that  sodium  metaphosphate  is  used  in  the  same  way* 
With  it  the  oxides  form  salts  of  phosphoric  acid. 

Most  of  the  boric  acid  obtained  from  Tuscany  is  used 
in  the  manufacture  of  borax,  a  salt  which  finds  extensive 
application. 

Salts  of  Boron.  —  Although  the  most  characteristic  com- 
pounds of  boron  are  those  in  which  it  acts  as  an  acid- 
forming  element,  it  forms  some  compounds  in  which  its 
power  as  a  base-former  is  shown.  Thus,  with  concen- 
trated sulphuric  acid  the  trioxide  forms  a  compound 
which  appears  to  be  pyrosulphuric  acid,  H2S2O7,  in  which 
one  hydrogen  is  replaced  by  the  group  BO,  which 
is  analogous  to  antimonyl,  SbO,  and  bismuthyl,  BiO. 
It  has  the  composition  (BO)HS2O7.  Further,  when  con- 
centrated phosphoric  acid  acts  upon  crystallized  boric 
acid,  boron  phosphate,  BPO4,  is  formed.  This  compound 
is  characterized  by  great  stability.  Concentrated  acids, 
for  example,  do  not  decompose  itt  It  also  forms  a  salt 
which  appears  to  be  analogous  to  tartar  emetic,  which, 
as  has  been  pointed  out,  is  probably  antimonyl  potas- 

sium tartrate,  C4H4O6  -j  ^     .    This  is  the  salt  represented 


(  "RO 

by  the  formula  C4H4O6  \  -g-    ,  which  may  be  called  boryl 

potassium  tartrate. 

Nitrogen  Boride,  BN.  —  This  Compound  has  been  re- 
ferred to  in  connection  with  the  preparation  of  boron. 
It  is  easily  obtained  by  igniting  a  mixture  of  dehydrated 
borax  and  ammonium  chloride.  It  forms  a  white  pow- 
der, which  is  insoluble  in  water,  and  is  characterized  by 
great  stability.  At  red  heat  it  is  decomposed  by  water 
vapor  into  ammonia  and  boric  acid  : 

2BN  +  6H30  =  2B(OH)3  +  2NH,. 


IJJ.S 


CHAPTER  XIX. 

CARBON  (C,  At.  Wt.  11.92)  AND  ITS  SIMPLER  COMPOUNDS 
WITH  HYDROGEN  AND  CHLORINE. 

Introductory. — Carbon  bears  to  Family  IV  relations 
similar  to  those  which  nitrogen,  oxygen,  and  fluorine 
bear  to  Families  V,  VI,  and  VII.  Towards  hydrogen, 
as  well  as  towards  chlorine  and  oxygen,  carbon  is  quadri- 
valent, and  towards  oxygen  it  is  also  bivalent.  In  this 
family  the  maximum  oxygen-valence  coincides  with  the 
hydrogen- valence,  while,  as  has  been  seen,  in  Families  V, 
VI,  and  VII,  the  oxygen- valence  is  higher  than  the  hy- 
drogen-valence, the  difference  becoming  greater  from 
Family  V  to  VII.  While  the  higher  oxygen  compounds 
of  Family  IV  are  acidic,  forming  acids  which  are  derived 
from  the  normal  acid,  E(OH)4,  the  lower  oxides  are  not 
generally  acid.  The  hydrogen  compounds  of  the  general 
formula  MH4,  of  which  there  are  but  two,  those  of  car- 
bon and  silicon,  have  neither  acid  nor  basic  properties. 
Carbon  is  distinguished  by  the  large  number  of  the 
compounds  into  which  it  enters,  all  of  which  are  more  or 
less  closely  related  to  a  comparatively  small  number  of 
fundamental  forms.  Silicon  also  forms  a  large  number 
of  compounds,  as  we  shall  see  ;  but  these  are  of  a  differ- 
ent kind  from  those  obtained  from  carbon. 

Occurrence  of  Carbon. — In  general,  substances  which 
are  obtained  from  the  vegetable  or  animal  kingdom  blacken 
when  heated  to  a  sufficiently  high  temperature,  and  after- 
wards, if  they  are  heated  in  the  air,  they  burn  up,  as  we 
say.  When  we  consider  the  great  variety  of  substances 
found  in  living  things,  it  certainly  appears  remarkable 
that  nearly  all  have  this  property  in  common.  It  is  due 
to  the  fact  that  nearly  all  animal  and  vegetable  substances 
contain  the  element  carbon.  When  they  are  heated  the 

(357) 


358  INORGANIC  CHEMISTRY. 

other  elements  present  are  first  driven  off  in  various 
forms  of  combination,  while  the  carbon  is  the  last  to  go. 
Hydrogen  and  oxygen  pass  off  as  water ;  hydrogen  and 
nitrogen  as  ammonia  ;  and  much  of  the  carbon  also  passes 
off  in  combination  with  hydrogen,  with  hydrogen  and 
oxygen,  and  with  nitrogen  and  hydrogen.  If  the  heat- 
ing is  carried  on  in  the  air,  the  carbon  finally  combines 
with  oxygen  to  form  a  colorless  gas — it  burns  up.  Car- 
bon is  the  central  element  of  organic  nature.  There  is 
not  a  living  thing,  from  the  minutest  microscopic  animal 
to  the  mammoth,  from  the  moss  to  the  giant  tree,  which 
does  not  contain  this  element  as  an  essential  constituent. 
The  number  of  the  compounds  which  it  forms  is  almost 
infinite,  and  they  present  such  peculiarities  that  they  are 
commonly  treated  of  under  a  separate  head,  "  Organic 
Chemistry."  There  is  no  good  reason  for  this,  except  the 
large  number  of  the  compounds.  For  our  present  pur- 
pose it  will  suffice  to  consider  the  chemistry  of  the  ele- 
ment itself,  and  of  a  few  of  its  more  important  simple 
compounds. 

From  what  has  already  been  said,  it  will  be  seen  that 
the  principal  form  in  which  carbon  occurs  in  nature  is 
in  combination  with  other  elements.  It  occurs  not  only 
in  living  things,  but  in  their  fossil  remains,  as  in  coal. 
Coal-oil,  or  petroleum,  the  formation  of  which  is  perhaps 
due  to  the  action  of  water  on  metallic  carbides,  consists 
of  a  large  number  of  compounds  which  contain  only  car- 
bon  and  hydrogen.  Most  products  of  plant-life  contain 
the  elements  carbon,  hydrogen,  and  oxygen.  Among  the 
more  common  of  these  products  may  be  mentioned  sugar, 
starch,  and  cellulose.  Most  products  of  animal  life  con- 
tain carbon,  hydrogen,  oxygen,  and  nitrogen.  Among 
them  may  be  mentioned  albumen,  fibrin,  casein,  etc. 
Carbon  occurs  in  the  air  in  the  form  of  carbon  dioxide. 
It  also  occurs  in  the  form  of  salts  of  carbonic  acid ;  the 
carbonates,  which  are  very  widely  distributed,  forming 
whole  mountain  ranges.  Limestone,  marble,  and  chalk 
are  varieties  of  calcium  carbonate. 

Uncombined,  the  element  occurs  pure  in  two  very  dif- 


DIAMOND—GRAPHITE.  359 

ferent  forms  in  nature :  (1)  As  diamond ;  and  (2)  as 
graphite,  or  plumbago. 

Diamond. — The  diamond  occurs  in  but  few  places 
on  the  earth,  and  practically  nothing  is  known  as  to  the 
conditions  which  gave  rise  to  its  formation.  The  cele- 
brated diamond  beds  are  in  India,  Borneo,  Brazil,  and 
South  Africa.  When  found,  diamonds  are  generally 
covered  with  an  opaque  layer,  which  must  be  removed 
before  its  beautiful  properties  are  apparent.  The  crys- 
tals are  sometimes  regular  octahedrons,  though  usually 
they  are  somewhat  more  complicated,  and  the  faces  are 
frequently  curved.  It  is  the  hardest  substance  known. 
For  use  as  a  gem  it  must  be  cut  and  polished. 
The  object  in  view  is  to  bring  out  as  strikingly  as  possi- 
ble its  brilliancy  by  exposing  the  faces  favorably  to 
the  action  of  the'  light.  If  heated  to  a  very  high  tem- 
perature without  access  of  air,  it  swells  up  and  is 
converted  into  a  black  mass  resembling  coke.  The 
change  takes  place  without  loss  in  weight.  Heated  to  a 
high  temperature  in  oxygen,  it  burns  up,  yielding  only 
carbon  dioxide.  It  is  insoluble  in  all  known  liquids  at 
ordinary  temperatures.  It  dissolves,  however,  in  molten 
cast  iron  and  in  some  other  molten  metals.  Small 
diamonds  have  recently  (1897)  been  made  by  Moissan  by 
dissolving  carbon  in  cast  iron  with  the  aid  of  an  electric 
furnace,  and  suddenly  cooling  the  mass.  Under  these 
conditions  the  carbon  is  under  great  pressure. 

Graphite. — Graphite,  or  plumbago,  is  found  in  nature 
in  large  quantities.  Sometimes  it  is  crystallized,  but  in 
forms  entirely  different  from  those  assumed  by  the  dia- 
mond. It  can  be  prepared  artificially  by  dissolving  char- 
coal in  molten  iron,  from  which  solution,  on  cooling,  it  is 
partly  deposited  as  graphite.  It  has  a  grayish-black 
color  and  a  metallic  lustre.  It  is  quite  soft,  leaving  a 
leaden-gray  mark  on  paper  when  drawn  across  it,  and  it 
is  hence  used  in  the  manufacture  of  so-called  lead  pencils. 
It  is  sometimes  called  black-lead.  When  heated  without 
access  of  air  it  remains  unchanged.  Heated  to  a  very  high 
temperature  in  the  air,  or  in  oxygen,  it  burns  up,  forming 


360  INORGANIC  CHEMISTRY. 

only  carbon  dioxide.  Like  the  diamond,  it  is  insoluble 
in  all  known  liquids  at  ordinary  temperatures. 

Amorphous  Carbon. — All  forms  of  carbon  which  are 
not  diamond,  nor  graphite,  are  included  under  the  name 
amorphous  carbon.  The  name  signifies  simply  that  it  is 
not  crystallized.  The  most  common  form  of  amorphous 
carbon  is  ordinary  charcoal. 

Charcoal  is  that  form  of  carbon  made  by  the  charring 
process,  which  consists  simply  in  heating  wood  without 
a  sufficient  supply  of  air  to  effect  complete  combus- 
tion. The  substance  almost  exclusively  used  in  the 
manufacture  of  charcoal  is  wood.  As  has  already  been 
stated,  wood  is  made  up  of  a  large  number  of  substances, 
nearly  all  of  which,  however,  consist  of  the  three  elements 
carbon,  hydrogen,  and  oxygen.  One  of  the  chief  con- 
stituents of  all  kinds  of  wood  is  cellulose.  Now,  when  we 
set  fire  to  a  piece  of  wood, — that  is  to  say,  when  we  heat 
it  up  to  the  temperature  at  which  oxygen  begins  to  act 
on  it, — it  burns,  if  air  is  present.  Under  ordinary  cir- 
cumstances the  chemical  changes  which  take  place  are 
complex;. but  if  care  is  taken,  the  combustion  can  be 
made  complete,  when  all  the  carbon  is  converted  into  car- 
bon dioxide,  and  all  the  hydrogen  into  water.  If,  on  the 
other  hand,  the  air  is  prevented  from  coming  in  contact 
with  the  wood,  as  by  heating  it  in  a  closed  vessel,  or  if  it 
is  prevented  from  coming  in  contact  with  it  sufficiently 
to  effect  complete  combustion,  the  hydrogen  is  given 
off  partly  as  water  and  partly  in  the  form  of  volatile 
products  containing  carbon  and  oxygen,  as  wood  spirits 
or  methyl  alcohol,  pyroligneous  or  acetic  acid,  acetone,  etc. 
The  carbon,  however,  is  for  the  most  part  left  behind 
as  charcoal,  as  there  is  not  enough  oxygen  to  convert 
it  into  carbon  dioxide.  Such  a  process  as  that  just  de- 
scribed, when  carried  on  in  closed  vessels,  is  known  as 
destructive  distillation  or  dry  distillation.  It  is  also  known 
as  the  charring  process.  It  is  a  complex  example  of  a 
kind  of  change  which  we  have  already  had  to  deal  with. 
Whenever  chemical  compounds  are  heated  the  constitu- 
ents tend  to  arrange  themselves  in  forms  which  are  stable 
at  the  higher  temperature.  Sulphites  become  sulphates ; 


CHARCOAL.  361 

phosphites  become  phosphates;  chlorates  become  per- 
chlorates ;  ammonium  salts  break  down  into  the  acids  and 
ammonia ;  ammonium  nitrite  is  decomposed  into  nitro- 
gen and  water ;  ammonium  nitrate  yields  nitrous  oxide 
and  water ;  primary  phosphates  yield  metaphosphates ; 
secondary  phosphates  yield  pyrophosphates,  etc.,  etc. 
Carbon  compounds  are,  in  general,  more  sensitive  to 
the  influence  of  heat  than  the  compounds  of  other 
elements,  and  all  are  decomposed  even  at  comparatively 
low  temperatures. 

The  above  statements  will  make  it  possible  to  under- 
stand the  working  of  a  charcoal-kiln.  This  consists  es- 
sentially of  a  pile  of  wood  arranged  to  leave  spaces  be- 
tween the  pieces,  and  covered  with  some  rough  material 
through  which  the  air  will  not  pass  easily,  as,  for  exam- 
ple, a  mixture  of  powdered  charcoal,  turf,  and  earth. 
Small  openings  are  left  in  this  covering,  so  that  after  the 
wood  is  kindled  it  will  continue  to  burn  slowly.  The 
process  is  sometimes  carried  on  in  structures  of  brick- 
work with  the  necessary  number  of  small  openings  in  the 
walls.  The  changes  above  mentioned  take  place,  the 
gases  or  volatile  substances  passing  out  of  the  top  of  the 
kiln,  and  appearing  as  a  dense  cloud.  In  due  time 
the  holes  through  which  the  air  gains  access  to  the 
wood,  thus  making  the  burning  possible,  are  stopped  up, 
and  the  burning  ceases.  Charcoal,  which  is  impure 
amorphous  carbon,  is  left  behind.  As  wood  always 
contains  some  incombustible  substances  in  small  quan- 
tity, these  are,  of  course,  found  in  the  charcoal.  When 
the  wood  or  charcoal  is  burned,  these  substances  remain 
behind  as  the  ash. 

Ordinary  charcoal  is  a  black,  comparatively  soft  sub- 
stance. It  burns  in  the  air,  though  not  easily,  unless 
the  gases  which  are  formed  are  constantly  removed  and 
fresh  air  is  supplied, — conditions  which  are  met  by  a 
good  draught,  or  by  blowing  upon  the  fire  with  a  bellows. 
It  burns  readily  in  oxygen.  The  product  of  the  com- 
bustion in  the  air  and  in  oxygen,  when  the  conditions  are 
favorable,  is  carbon  dioxide,  CO2.  In  the  air,  when  the 
draught  is  bad,  another  compound  of  carbon  and  oxygen, 


362  INORGANIC  CHEMISTRY. 

carbon  monoxide,  CO,  is  formed.  Heated  without  access 
of  air,  charcoal  remains  unchanged.  Charcoal  is  insolu- 
ble in  liquids  generally,  though  it  is  soluble  in  molten 
iron,  and  it  crystallizes  from  the  solution,  as  we  have 
seen,  in  the  form  of  graphite,  and  sometimes  as  diamond. 

Coke. — Besides  wood  charcoal,  there  are  other  forms 
of  amorphous  carbon,  which  are  manufactured  for  special 
purposes,  or  are  formed  in  processes  carried  on  for  the 
sake  of  other  products.  Coke  is  a  form  of  amorphous 
carbon  which  is  made  by  heating  ordinary  gas-coal  with- 
out access  of  air,  as  is  done  on  the  large  scale  in  the 
manufacture  of  illuminating  gas.  Coke  bears  to  coal 
much  the  same  relation  that  charcoal  bears  to  wood. 

Lamp-black  is  a  very  finely  divided  form  of  charcoal 
which  is  deposited  on  cold  objects  placed  in  the  flames 
of  burning  oils.  The  oils  consist  almost  exclusively  of 
carbon  and  hydrogen.  When  burned  in  the  air  they 
yield  carbon  dioxide  and  water.  If  the  flame  is  cooled 
down  by  any  means,  or  if  the  supply  of  air  is  partly  cut 
off,  the  carbon  is  not  completely  burned,  the  flame 
"smokes,"  as  we  say,  and  deposits  soot.  This  process 
is  chemically  analogous  to  the  deposit  of  metallic  arsenic 
from  a  flame  of  arsine,  to  the  deposit  of  sulphur  from 
a  flame  of  hydrogen  sulphide,  and  that  of  phosphorus 
from  a  flame  of  phosphine,  when  these  gases  are  burned 
in  a  supply  of  air  insufficient  to  effect  complete  combus- 
tion of  both  constituents.  The  soot  obtained  from  the 
flames  of  burning  oils  is  made  up  largely  of  fine  particles 
of  carbon,  though  some  of  the  unchanged  oils  are  con- 
tained in  it.  It  is  used  in  the  manufacture  of  printing- 
ink.  As  carbon  is  acted  upon  directly  by  very  few  sub- 
stances, and  is  not  soluble,  it  is  almost  impossible  to 
destroy  the  color  of  printing-ink  without  destroying  the 
material  upon  which  it  is  impressed. 

Bone-Hack,  or  Animal  Charcoal,  is  a  form  of  amorphous 
carbon  which  is  made  by  charring  bones.  Bones  consist 
of  about  one  third  organic  matter  and  two  thirds  incom- 
bustible matter,  mostly  calcium  phosphate.  When 
charred,  the  organic  matter  undergoes  the  changes 
briefly  described  under  the  head  of  Charcoal,  while  the 


CHARCOAL.  363 

incombustible  constituents  remain  unchanged.  As  the 
organic  matter  is  distributed  through  the  substance  of 
the  bones  the  charcoal  obtained  in  this  way  is  in  a  very 
fine  state  of  division,  but  it  is  mixed  with  several  times 
its  own  weight  of  mineral  matter.  In  order  to  remove 
the  latter  the  bone-black  must  be  treated  with  hydro- 
chloric acid,  and  afterwards  thoroughly  washed  with 
water.  An  efficient  variety  of  animal  charcoal  is  made, 
further,  by  mixing  blood  with  sodium  carbonate,  char- 
ring, and  afterwards  dissolving  out  the  sodium  carbon- 
ate with  water. 

Bone-black  and  wood-charcoal  are  very  porous,  and 
have  the  power  to  absorb  gases.  When  placed  in  air 
containing  bad-smelling  gases  these  are  absorbed,  and 
the  air  is  thus  to  some  extent  purified.  When  water  con- 
taining disagreeable  substances  is  treated  with  charcoal, 
these  are  wholly  or  partly  absorbed,  and  the  water  is  im- 
proved. Charcoal-filters  are  therefore  extensively  used. 
A  charcoal-filter  to  be  efficient  should  be  of  good  size, 
and  from  time  to  time  the  charcoal  should  be  taken  out 
and  renewed.  The  small  filters  which  are  screwed  into 
faucets  are  of  little  value,  as  the  charcoal  soon  becomes 
charged  with  the  objectionable  material  which  is  pres- 
ent in  the  water,  and  is  then  a  source  of  contamination 
rather  than  a  means  of  purification.  The  power  of 
charcoal  to  absorb  gases  depends  upon  its  porosity. 
That  from  some  varieties  of  wood  is  more  porous  than 
that  from  other  varieties.  Box-wood  charcoal  has 
been  shown  to  absorb  90  times  its  own  volume  of  am- 
monia gas,  35  times  its  volume  of  carbon  dioxide,  and 
nearly  twice  its  volume  of  hydrogen.  Charcoal  from 
cocoa-nut  wood  absorbs  172  times  its  volume  of  am- 
monia, and  68  times  its  volume  of  carbon  dioxide. 

Some  coloring  matters  can  be  removed  from  liquids 
by  passing  the  liquids  through  bone-black  filters.  On  the 
large  scale,  this  fact  is  taken  advantage  of  in  the  refining 
of  sugar.  The  solution  of  sugar  first  obtained  from  the 
cane  or  beet  is  highly  colored ;  and,  if  it  were  evapo- 
rated, the  sugar  deposited  from  it  would  be  dark-colored. 
If,  however,  the  solution  is  first  passed  through  bone- 


364  INORGANIC  CHEMISTRY. 

black  filters,  the  color  is  removed,  and  now,  on  evaporat- 
ing, white  sugar  is  deposited.  In  the  laboratory  con- 
stant  use  is  made  of  this  method  for  decolorizing  liquids. 
The  action  can  easily  be  shown  by  adding  a  little  bone- 
black  to  a  solution  containing  some  litmus  or  indigo. 
If  the  solution  is  digested  for  a  short  time  with  the 
bone-black,  and  then  passed  through  a  filter,  it  will  be 
found  that  the  coloring  matter  is  removed. 

Charcoal  does  not  undergo  decay  in  the  air  or  under 
water  nearly  as  readily  as  wood.  That  is  another  way 
of  stating  the  chemical  fact  that  the  substances  of  which 
wood  is  made  up  are  more  susceptible  to  the  action  of 
other  chemical  substances  than  charcoal  is.  We  have 
one  good  illustration  of  this,  indeed,  in  the  relative  ease 
with  which  charcoal  and  wood  burn  in  the  air.  Piles 
which  are  driven  below  the  surface  of  water  are  some- 
times charred  to  protect  them  from  the  action  of  those 
substances  which  cause  decay. 

Coal. — Under  this  head  are  included  a  great  many 
kinds  of  impure  amorphous  carbon  which  occur  in  na- 
ture. Although  we  might  distinguish  between  an  almost 
infinite  number  of  kinds  of  coal,  for  ordinary  purposes 
they  are  divided  into  hard  and  soft  coals,  or  anthracite 
and  bituminous  coals.  Then  there  are  substances  more 
nearly  allied  to  wood  called  lignite,,  and  those  which  rep- 
resent a  very  early  stage  in  the  process  of  coal-forma- 
tion, viz.,  peat.  A  close  examination  01  all  these  varie- 
ties has  shown  that  they  have  been  formed  by  the 
gradual  decomposition  of  vegetable  matter  in  an  insuf- 
ficient supply  of  air.  The  process  has  been  going  on 
for  ages.  Sometimes  the  substances  have,  at  the  same 
time,  been  subjected  to  great  pressure;  ac  can  be  seen 
from  the  position  in  which  they  occur  in  the  earth.  The 
products  in  the  earlier  stages  of  the  coal-forming  pro- 
cess are  more  closely  allied  to  wood  than  those  in  the 
later  stages.  All  forms  of  coal  contain  other  substances 
in  addition  to  the  carbon.  The  soft  coals  are  particu-= 
larly  rich  in  other  substances.  When  heated  they  give 
off  a  mixture  of  gases  and  the  vapors  of  volatile  liquids. 
The  gases  are,  for  the  most  part,  useful  for  illuminating 


DIAMOND,  GRAPHITE,  AND  CHARCOAL.  365 

purposes.  The  liquids  form  a  black,  tarry  mass  known 
as  coal-tar,  from  which  many  valuable  compounds  of  car- 
bon are  obtained.  The  gases  are  passed  through  water 
for  the  purpose  of  removing  certain  impurities.  This 
water  absorbs  ammonia,  and  forms  the  ammoniacal  liquor 
of  the  gas-works,  which,  as  has  been  stated,  is  the  prin- 
cipal source  of  ammonia. 

Diamond,  Graphite,  and  Charcoal  are  Different  Forms  of 
the  Element  Carbon. — We  have  seen  that  oxygen  presents 
itself  in  two  forms — ordinary  oxygen  and  ozone.  Ozone 
is  made  from  oxygen,  and  oxygen  from  ozone,  without  any 
increase  or  decrease  in  weight ;  and  the  compounds  ob- 
tained by  the  combination  of  other  elements  with  oxygen 
are  identical  with  those  obtained  by  the  combination  of 
the  same  elements  with  ozone.  So  also  there  are  several 
varieties  of  sulphur,  two  of  which  crystallize  in'  different 
forms.  There  are,  further,  three  or  four  different  modi- 
fications of  the  element  phosphorus,  and  these  differ 
from  one  another  in  a  very  marked  way.  The  explana- 
tion of  the  difference  between  oxygen  and  ozone  is  that 
the  molecule  of  the  former  is  made  up  of  two  atoms, 
while  that  of  ozone  is  made  up  of  three,  which  are  in 
a  state  of  unstable  equilibrium.  This  explanation  is 
reached  through  a  study  of  the  specific  gravity  of  the 
two  gases.  At  present  no  satisfactory  explanation  can 
be  given  of  the  difference  between  the  varieties  of  phos- 
phorus and  between  the  varieties  of  sulphur.  It  will 
probably  be  shown  to  be  due  to  the  way  in  which  the 
atoms  are  grouped  together  in  the  molecules,  and  also 
to  the  way  in  which  the  molecules  are  grouped  together 
to  form  the  masses.  Carbon,  as  we  have  seen,  occurs  in 
three  distinct  forms0  It  is  difficult  to  conceive  that  the 
black,  porous  charcoal,  and  the  dull,  gray,  soft  graphite 
are  chemically  identical  with  the  hard,  transparent, 
brilliant  diamond.  Yet  this  is  undoubtedly  the  case,  as 
can  be  shown  by  a  very  simple  experiment.  Each  of 
the  substances  when  burned  in  oxygen  yields  carbon 
dioxide.  Now?  the  composition  of  carbon  dioxide  is 
known,  so  that,  if  the  weight  of  the  carbon  dioxide 
formed  in  a  given  experiment  is  known,  the  weight  of  the 


366  INORG'ANIC  CHEMISTRY. 

carbon  in  it  is  also  known.  When  a  gram  of  pure  char- 
coal is  burned  it  yields  3f  grams  carbon  dioxide,  and  in 
this  quantity  of  carbon  dioxide  there  is  contained  ex- 
actly one  gram  of  carbon.  Further,  when  a  gram  of 
graphite  is  burned  the  same  weight  (3f  grams)  of  carbon 
dioxide  is  obtained  as  in  the  case  of  charcoal ;  and  the 
same  thing  is  true  of  diamond.  It  follows  from  these 
facts  that  the  three  forms  of  matter  known  as  char- 
coal, graphite,  and  diamond  consist  only  of  the  element 
carbon.  The  explanation  of  the  diffei  ^nce  is  not  known, 
but,  as  in  the  cases  of  phosphorus  an  \  sulphur,  it  will 
probably  be  found  to  be  in  the  differeii  ways  in  which 
the  atoms  are  arranged  in  the  molecule,  and  the  mole- 
cules in  the  masses. 

Notwithstanding  the  marked  differences  in  their  ap- 
pearance and  in  many  of  their  physical  properties,  the 
three  forms  of  carbon  have,  as  we  have  seen,  some  prop- 
erties in  common.  They  are  insoluble  in  all  known 
liquids  at  ordinary  temperatures.  They  are  tasteless, 
inodorous,  and  infusible.  When  heated  without  access 
of  air  they  remain  unchanged,  unless  the  temperature  is 
very  high,  when  the  diamond  swells  up  and  is  converted 
into  a  mass  resembling  coke — a  change  which  is  con- 
nected with  a  rearrangement  of  the  particles  in  an  irreg- 
ular way,  so  that  the  substances  cease  to  be  crystalline, 
or  become  amorphous. 

Chemical  Conduct  of  Carbon. — At  ordinary  tempera- 
tures carbon  is  an  inactive  element.  If  it  is  left  in  contact 
with  any  one  of  the  elements,  no  chemical  change  takes 
place.  It  will  not  combine  with  any  of  them  unless  the  tem- 
perature is  raised.  At  higher  temperatures,  however,  it 
combines  with  several  of  them  with  great  ease,  especially 
with  oxygen.  Under  proper  conditions  it  combines  also 
with  nitrogen,  with  hydrogen,  with  sulphur,  and  with 
many  other  elements.  It  combines  with  oxygen  either 
directly,  as  when  it  burns  in  the  air  or  in  oxygen ;  or  it 
abstracts  oxygen  from  some  of  the  oxides.  The  direct 
combination  of  oxygen  and  carbon  has  already  been  seen 
in  the  burning  of  charcoal  in  oxygen,  and  is  familiar  to 
every  one  in  the  fire  in  a  charcoal  furnace.  That  car- 


CHEMICAL  CONDUCT  OF  CARBON.  367 

ton  dioxide  is  the  product  formed  can  be  shown  by 
passing  the  gas  through  lime-water  or  baryta-water, 
when  insoluble  calcium  or  barium  carbonate  will  be 
thrown  down.  The  reason  why  lime-water  or  baryta- 
water  is  used  is  simply  that  an  insoluble  compound  is 
formed,  and  this  can  be  seen,  and  it  can  be  separated 
from  the  liquid  and  examined.  The  reaction  which 
takes  place  is  represented  thus : 

Ca(OH)3  +  CO2  =  CaCO3  +  H2O; 

Calcium  Carbon         Calcium 

hydroxide        dioxide       carbonate 

Ba(OH)2  +  C02  =  BaC03  +H2O. 

Barium  Carbon          Barium 

hydroxide        dioxide       carbonate 

No  other  common  gas  acts  in  this  way  on  these  solu- 
tions. Hence,  when,  under  ordinary  circumstances,  a 
gas  is  passed  into  lime-water  and  an  insoluble  compound 
is  formed,  we  may  conclude  that  the  gas  is  carbon  diox- 
ide, though  this  conclusion  may  require  further  proof. 

The  abstraction  of  oxygen  from  compounds  by  means 
of  carbon  may  be  illustrated  in  a  number  of  ways.  Thus, 
when  powdered  copper  oxide,  CuO,  is  mixed  with  pow- 
dered charcoal,  and  the  mixture  heated  in  a  tube,  car- 
bon dioxide  is  given  off,  and  can  be  detected  as  in  the 
last  experiment  mentioned.  Copper  is  left  behind,  and, 
if  the  proportions  are  properly  selected,  all  the  carbon 
will  pass  off  as  carbon  dioxide,  and  only  the  copper  be 
left  behind : 

2CuO  +  C  =  2Cu  +  CO,. 

In  a  similar  way,  arsenious  oxide,  As^O3,  gives  up  its 
oxygen  to  carbon.  This  fact  furnishes  indeed  a  delicate 
method  for  the  detection  of  the  substance.  If  a  little 
is  placed  in  th'3  bottom  of  a  small  tube,  and  above  it  a 
small  piece  of  charcoal,  then  when  heat  is  applied  the 
arsenious  oxide  sublimes,  and  as  its  vapor  passes  the 
heated  charcoal  the  oxygen  is  abstracted,  and  the  ele- 
ment arsenic,  being  also  somewhat  volatile,  is  deposited 


368  INORGANIC  CHEMISTRY. 

just  above  the  charcoal  in  the  form  of  a  lustrous  mirror 
on  the  walls  of  the  tube.     The  reaction  is 


As  has  already  been  explained,  the  abstraction  of  oxy- 
gen from  a  compound  is  known  as  reduction.  Hence, 
carbon  is  called  a  reducing  agent.  It  is  indeed  the  re- 
ducing agent  which  is  most  extensively  used  in  the  arts. 
Its  chief  use  is  in  extracting  metals  from  their  ores, 
which  are  the  forms  in  which  they  occur  in  nature. 
Thus,  iron  does  not  occur  in  nature  as  iron,  but  in  com- 
bination with  other  elements,  especially  with  oxygen. 
In  order  to  get  the  metal  the  ore  must  be  reduced,  or, 
in  other  words,  the  oxygen  must  be  extracted.  This  is 
invariably  accomplished  by  heating  it  with  some  form 
of  carbon,  either  coke  or  charcoal. 

The  elements  chlorine,  oxygen,  nitrogen,  and  hydro- 
gen being  gases,  and  the  products  formed  when  the  first 
three  combine  with  hydrogen  being  also  gaseous  or  con- 
vertible into  vapor,  it  is  a  comparatively  easy  matter  to 
study  the  relations  between  the  volumes  of  the  combin- 
ing gases  and  the  volumes  of  the  products  formed.  It  is, 
however,  impossible  to  determine  the  ratio  between  the 
volume  of  carbon  gas  and  that  of  other  gases  with  which 
it  combines. 

Compounds  of  Carbon  with  Hydrogen,  or  Hydrocarbons. 
Conditions  under  which  Hydrocarbons  are  Formed.  —  When 
the  carbon  pencils  connected  with  a  powerful  battery, 
as  in  the  production  of  the  electric  arc-light,  are  sur- 
rounded by  an  atmosphere  of  hydrogen  the  two  elements 
combine  to  some  extent  to  form  the  compound  acetylene, 
C,Ha.  When  organic  matter  undergoes  decomposition 
without  free  access  of  air,  as  for  example  under  water, 
the  carbon  compounds  are  reduced  to  the  final  product 
known  as  marsh-gas  or  methane,  CH4,  just  as  compounds 
containing  nitrogen  yield  ammonia.  Compounds  of 
various  metals  with  carbon,  known  as  carbides,  yield 
hydrocarbons,  such  as  acetylene  and  marsh-gas,  when 
treated  with  water.  The  compounds  which  make  up 
petroleum  are  hydrocarbons  which  have  probably  been 


NUMBER  OF  HYDROCARBONS.  369 

formed,  either  by  the  action  of  water  on  metallic  carbides 
in  the  interior  of  the  earth,  or  by  decomposition  of  organic 
matter  without  free  access  of  air.  Finally,  when  wood 
or  coal  is  heated,  hydrocarbons  are  given  off,  and  com- 
pounds of  this  kind  are  therefore  contained  in  coal-gas. 

Number  of  Hydrocarbons. — The  number  of  hydrocar- 
bons is  very  great,  and  new  ones  are  constantly  being 
made.  The  fact  that  carbon  is  distinguished  for  the 
large  number  of  its  compounds  has  already  been  men- 
tioned. The  simplest  of  these  are  the  hydrocarbons. 
It  is  safe  to  say  that  there  are  as  many  as  two  hun- 
dred hydrocarbons  known.  Fortunately,  however,  most 
of  these  have  been  found  to  bear  comparatively  simple 
relations  to  one  another,  and  therefore,  though  the 
number  is  large  and  the  variety  great,  their  study  is  not 
as  difficult  as  one  would  be  inclined  to  think. 

Petroleum  is  an  oily  liquid  found  in  many  places  in  the 
earth  in  large  quantity,  and  escapes  when  a  cavity  in 
which  it  is  contained  is  punctured.  In  the  earth  it 
contains  both  gases  and  liquids.  When  it  is  brought 
into  the  air  the  gases  and  the  liquids  which  are  vola- 
tile at  the  ordinary  temperature  are  given  off.  There 
are  several  gases  given  off,  and  a  large  number  of 
liquids  left  behind.  The  simplest  gas  corresponds  to 
the  formula  CH4,  the  next  to  C2H6,  the  next  to  C3H8,  the 
next  to  C4H10.  An  examination  of  the  liquid  has  shown 
that  it  contains  other  hydrocarbons  of  the  formulas 
C5H12,  C6H14,  C7H16,  C8H18,  etc.  It  will  be  seen  that  these 
compounds  bear  a  simple  relation  to  one  another,  as  far 
as  composition  is  concerned.  Arranging  them  in  a  series, 
tne  first  eight  members  are 

CH4,    Methane,  or  Marsh-gas  ; 

C2H6,   Ethane; 

C3H8,  Propane ; 

C4H10,  Butane ; 

C6H12,  Pentane ; 

C6H14,  Hexane ; 

C7H16,  Heptane ; 

C8H18,  Octane. 


370  INORGANIC  CHEMISTRY. 

Homology,  Homologous  Series. — In  the  above  series  the 
first  member  differs  from  the  second  by  CH2 ;  and  there 
is  also  this  same  difference  between  the  second  and 
third,  the  third  and  fourth,  and  in  general  between  any 
two  consecutive  members  in  the  series.  This  relation  is 
known  as  homology,  and  such  a  series  is  known  as  an 
homologous  series.  Carbon  is  distinguished  from  all 
other  elements  by  its  power  to  form  homologous  series. 
Other  elements  form  similar  series,  but  the  homology  is 
not  of  the  same  kind  as  that  which  is  met  with  among 
carbon  compounds.  Thus  the  series  of  chlorine  acids, 
and  the  similar  series  of  acids  of  nitrogen  and  sulphur, 
are  homologous  series,  in  which  the  constant  difference 
between  any  two  consecutive  members  is  represented  by 
an  atom  of  oxygen  : 

HC1O  H2SO2  HNO 

HC102  H2S03  HNO2 

HC103  H2SO4  HN03. 
HC104 

These  series,  are,  however,  much  less  extensive  than  the 
homologous  series  of  compounds  of  carbon. 

Cause  of  the  Homology  among  Compounds  of  Carbon. — 
The  explanation  of  the  homology  observed  between  the 
compounds  of  carbon  is  founded  on  the  view  that  carbon 
is  quadrivalent,  and  that  it  has  the  power  to  unite  with 
itself  in  chains.  The  quadrivalence  is  shown  in  the  com- 
pounds CH4,  CC14,  CHC13,  CO2,  etc.  When  marsh-gas, 
CH4,  is  treated  with  chlorine  the  first  product  of  the  ac- 
tion is  chlor-methane,  CH3C1,  which  according  to  the 
prevailing  views  is  marsh-gas  in  whose  molecule  one 
atom  of  hydrogen  has  been  replaced  by  chlorine.  The 

H 

structure  of  this  compound  is  represented  thus,  H-C-C1, 

*      H 
H 
if  that  of  marsh-gas  is  represented  in  this  way,  H-C-H. 


OTHER  SERIES  OF  HYDROCARBONS.  371 

Now,  when  chlor-methane  is  treated  with  sodium  the 
chlorine  is  extracted,  and  a  compound  of  the  formula 
CaH6  is  formed : 

2CH3C1  +  2Na  =  CaHfl  +  2NaCl. 

It  appears  that,  the  chlorine  being  extracted  from  the 
compound,  the  residues  of  the  composition  CH3  unite 
in  pairs  to  form  the  compound  C2H6,  which  is  ethane,  or 
the  second  member  of  the  series  of  hydrocarbons  above 
given.  The  simplest  explanation  of  the  facts  stated  is 
that  the  carbon  atoms  unite  by  means  of  the  bonds  or 
valences  left  free  when  the  chlorine  is  extracted.  The 
residues  after  the  extraction  of  the  chlorine  may  be  rep- 

H 

resented  by  the  formula  H-C-;  when    two    of    these 

i 

unite,  the  resulting  compound  will  have  the  structure 

H  H 

represented  by  the   formula  H-C-C-H.     In  a  similar 

ii  ii 

way  the  relation  between  the  other  members  of  the 
series  can  be  explained,  and  the  explanation  is  in  perfect 
accordance  with  a  large  number  of  facts. 

Other  Series  of  Hydrocarbons.—  Besides  the  series  above 
mentioned,  which,  as  its  simplest  member  is  marsh-gas, 
is  known  as  the  marsh-gas  series,  there  are  other  homolo- 
gous series  of  hydrocarbons.  There  is  one  beginning 
with  ethylene,  or  olefiant  gas,  C2H4,  examples  of  which 
are 

Ethylene,    C2H4 ; 

Propylene,  C3H6 ; 

Butylene,    C4H8. 

There  is  another  beginning  with  acetylene,  CaHa,  ex- 
amples of  which  are 

Acetylene,  C2Ha; 
Allylene,     C3H4. 


372  INORGANIC  CHEMISTRY. 

Another  series  begins  with  benzene,  C6Ha.     Some  of 
the  members  of  this  series  are 


Benzene,  C6H 
Toluene,  C,H 
Xylene,  C8H 


These  are  the  hydrocarbons  which  are  obtained  from 
coal-tar. 

The  relations  between  these  hydrocarbons  and  those 
of  the  marsh-gas  series  have  been  extensively  studied, 
and  a  great  deal  of  light  has  been  thrown  upon  the  sub- 
ject. It  would,  however,  lead  too  far  to  take  up  this 
subject  here.  A  word  may  be  said,  however,  in  regard 
to  the  relations  believed  to  exist  between  the  hydrocar- 
bons of  the  ethylene  and  the  acetylene  series,  and  those 
of  the  marsh-gas  series.  It  is  believed  that  in  ethylene 
the  two  carbon  atoms  are  united  in  a  different  way  from 
that  in  ethane.  The  condition,  whatever  it  may  be,  is 
thought  to  be  similar  to  that  which  exists  in  the  mole- 
cule of  a  compound  consisting  of  two  bivalent  atoms, 
as,  for  example,  calcium  oxide.  The  condition  is  called 
double  union,  and  it  is  represented  by  a  double  line  as 
in  the  formulas  Ca=O,  H2C=CH2,  etc.  Whatever  this 
condition  may  be,  it  carries  with  it  the  power  to  com- 
bine with  other  atoms.  Thus,  when  ethylene  is  treated 
with  chlorine  it  takes  up  two  atoms,  and  is  converted 
into  dichlorethane,  C2H4C12,  the  double  union  being  de- 
stroyed and  single  union  existing  in  the  resulting  com- 
pound as  in  ethane.  So  also,  when  calcium  oxide, 
Ca=O,  is  treated  with  water,  it  is  converted  into  the  hy- 
droxide, in  which  the  condition  of  double  union  does  not 
exist : 

Ca        OH        Ca-OH 

ii    +   i      =i          ; 
O         H  OH 

CH2        Cl        CH2C1 

II       +1=1 
CH2        Cl        CHaCl 

Similar  reasons  have  led  to  the  conclusion  that  in 
acetylene,  CaH2,  the  carbon  atoms  are  held  together  in 


MABSH-GAS,  METHANE,  FIRE-DAMP.  373 

a  -different  way  from  that  in  ethane,  and  that  in  ethy- 
lene.  This  is  believed  to  be  similar  to  the  kind  of  union 
which  exists  in  a  molecule  consisting  of  two  trivalent 
atoms,  as  in  the  compound  boron  nitride,  B=N,  a  con- 
dition which  is  called  triple  union  or  triple  linkage.  This 
view  is  expressed  by  the  formula  HC=CH.  Wherever 
this  condition  exists  we  find  the  power  to  take  up  four 
univalent  atoms,  the  compound  thus  becoming  saturated, 
as  we  say.  Thus,  acetylene  can  take  up  four  atoms  of 
chlorine,  or  two-  of  hydrogen  and  two  of  chlorine,  form- 
ing in  the  former  case  tetra-chlor-ethane,  and  in  the  latter 
di-chl or- ethane  : 

HteCH  +  4C1     =  C12HC-CHC12(C2H2C14) ; 
HC=CH  +  2HC1  =  C1H2C-CH2C1(C2H4C12). 

Marsh-gas,  Methane,  Fire-damp,  CH4. — Marsh-gas  is 
found  in  nature  in  petroleum,  and  is  given  off  when  the 
oil  is  taken  out  of  the  earth  and  the  pressure  removed. 
It  is  formed,  as  the  name  implies,  in  marshes,  as  the 
product  of  a  reducing  process.  Vegetable  matter  is  com- 
posed essentially  of  carbon,  hydrogen,  and  oxygen. 
When  it  undergoes  decomposition  in  the  air  in  a  free 
supply  of  oxygen,  the  final  products  formed  are  carbon 
dioxide  and  water.  When  the  decomposition  takes 
place  without  access  of  oxygen,  as  under  water,  marsh- 
gas,  which  is  a  reduction-product,  is  formed.  The  gas 
can  be  made  in  the  laboratory  by  passing  a  mixture  of 
hydrogen  sulphide  and  the  vapor  of  carbon  disulphide 
over  heated  copper,  and  also  by  the  action  of  water  on 
several  metallic  carbides,  more  especially  aluminium 
carbide,  C3A14 : 

C3A14  +  12H20  =  3CH4  +  4A1(OH),. 

The  gas  is  met  with  in  coal-mines,  and  is  known  to 
the  miners  as  Jire-damp,  damp  being  the  general  name 
applied  to  a  gas,  and  the  name  fire-damp  meaning  a  gas 
that  burns.  To  prepare  it  in  the  laboratory  it  is  most 
convenient  to  heat  a  mixture  of  sodium  acetate  and 


374  INORGANIC  CHEMISTRY. 

quick  -lime.  The  change  which  takes  place  will  be  most 
readily  understood  by  regarding  it  as  a  simple  decom- 
position of  acetic  acid.  Acetic  acid  has  the  formula 
C2H4O2.  When  heated  alone  it  boils,  and  does  not  suffer 
decomposition.  If  it  is  converted  into  a  salt,  and  heated 
in  the  presence  of  a  base,  it  breaks  down  into  marsh-gas 
and  carbon  dioxide  : 


The  carbon  dioxide,  which  with  bases  forms  salts,  does 
not  pass  off,  but  remains  behind  in  the  form  of  a  salt  of 
carbonic  acid. 

Marsh-gas  is  a  colorless,  transparent,  tasteless,  inodor- 
ous gas.  It  is  slightly  soluble  in  water,  and  burns, 
forming  carbon  dioxide  and  water.  When  mixed  with 
air  the  mixture  explodes  if  a  flame  or  spark  comes  in 
contact  with  it.  This  is  one  of  the  causes  of  the  explo- 
sions which  so  frequently  occur  in  coal-mines.  To  pre- 
vent these  explosions  a  special  lamp  was  invented  by  Sir 
Humphry  Davy,  which  is  known  as  the  safety  -lamp.  The 
simple  principles  involved  in  its  construction  will  be  ex- 
plained when  the  subject  of  flame  is  taken  up. 

Ethylene,  defiant  Gas,  C2H4.  —  This  liydrocarbon  is 
formed  by  heating  a  mixture  of  ordinary  alcohol  and 
concentrated  sulphuric  acid.  The  reaction  is  represented 
thus  : 

C.H.O  -  H.O  +  CQH, 

Alcohol  Ethylene 

Ethylene  is  a  colorless  gas,  which  can  be  condensed  to 
a  liquid.  It  burns  with  a  luminous  flame,  and  forms  an 
explosive  mixture  with  oxygen. 

Acetylene,  C2H2.  —  Acetylene  is  formed  when  a  current 
of  hydrogen  is  passed  between  carbon  poles,  which  are 
incandescent  in  consequence  of  the  passage  of  a  power- 
ful electric  current.  In  this  case  carbon  and  hydrogen 
combine  directly.  It  is  formed  also  when  the  flame  of 
an  ordinary  laboratory  gas-burner  (Bunsen  burner) 
"  strikes  back,"  or  burns  at  the  base  without  a  free  sup- 


SIMPLER  COMPOUNDS  OF  CARBON.  375 

ply  of  air.     It  is  most  easily  obtained  by  treating  certain 
metallic  carbides,  especially  calcium  carbide,  with  water  : 

C2Ca  +  H20  =  C2H8  +  CaO. 

Its  odor  is  unpleasant.  It  burns  with  a  luminous, 
smoky  flame. 

Simpler  Compounds  of  Carbon  with,  the  Members  of  the 
Chlorine  Group. — When  chlorine  acts  upon  a  hydrocar- 
bon it  generally  takes  the  place  of  the  hydrogen,  atom 
for  atom.  Thus,  when  it  acts  upon  marsh-gas,  the 
following  reactions  take  place ; 

CH4  +  C12  =  CH3C1  +  HC1 ; 
CH3C1  +  C12  =  CH2C12  +  HC1 ; 
CH2C12  +  01,  =  CHC13  +  HC1 ; 
CHC13  +  01,  =  CC14  +  HC1. 

The  four  products  are  known  respectively  as  mono-cUor- 
methane,  di-cJdor-methane,  tri-chlor-metharie,  and  tetra-chlor- 
methane,  or  carbon  tetrachloride.  By  treating  these  com- 
pounds with  nascent  hydrogen  they  can  all  be  converted 
back  again  into  marsh-gas.  The  fact  that  the  hydrogen 
in  marsh-gas  can  be  replaced  in  four  steps,  one  fourth 
of  the  hydrogen  being  replaced  at  each  step,  furnishes  a 
strong  confirmation  of  the  correctness  of  the  view  ex- 
pressed by  the  formula  CH4,  which  signifies  that  in  the 
molecule  of  marsh-gas  there  are  four  atoms  of  hydrogen. 
The  most  important  of  the  four  compounds  is  the  third, 
tri-chlor-methane  or  chloroform.  While  chloroform  can 
be  made  by  treating  marsh-gas  with  chlorine,  it  is  much 
more  easily  obtained  in  other  ways,  as,  for  example,  by 
treating  alcohol  or  acetone  with  bleaching-powder. 
Without  a  study  of  the  relations  which  exist  between 
several  classes  of  compounds  of  carbon  these  reactions 
cannot  well  be  explained,  and  their  study,  as  well  as  that 
of  chloroform,  had  better  be  postponed  until  the  subject 
of  Organic  Chemistry  is  taken  up  systematically.  Cor- 
responding to  chloroform  there  are  bromine  and  iodine 
derivatives,  known  as  bromoform,  CHBr3,  and  iodoform, 
CHI3. 


CHAPTER  XX. 


SIMPLER  COMPOUNDS  OF  CARBON  WITH  OXYGEN,  AND 
WITH  OXYGEN  AND  HYDROGEN. 

General.  —  The  final  product  of  oxidation  of  carbon  is 
carbon  dioxide,  and  the  final  product  of  reduction  is 
marsh-gas,  but  between  these  two  limits  there  are  a 
number  of  interesting  derivatives,  just  as  there  are  a 
number  of  compounds  of  sulphur  between  hydrogen  sul- 
phide and  sulphuric  acid  ;  a  number  of  compounds  of 
nitrogen  between  ammonia  and  nitric  acid  ;  and  a  number 
of  compounds  of  phosphorus  between  phosphine  and 
phosphoric  acid.  We  have  seen  that  in  the  cases  men- 
tioned two  kinds  of  change  are  brought  about  by  oxida- 
tion :  (1)  Owing  to  the  fact  that  the  valence  of  chlo- 
rine, sulphur,  nitrogen,  and  phosphorus  towards  oxygen 
is  greater  than  towards  hydrogen,  the  act  of  oxidation 
involves  the  addition  of  oxygen  to  the  element  ;  (2)  hy- 
droxyl  appears  to  take  the  place  of  the  hydrogen  atoms 
one  by  one.  In  the  case  of  carbon,  the  valence  towards 
hydrogen  and  oxygen  being  the  same,  only  the  latter 
kind  of  change  takes  place. 

Relations  between  the  Compounds  of  Carbon  with  Hy- 
drogen and  Oxygen.  —  In  order  that  the  relations  between 
the  simpler  compounds  of  carbon  with  hydrogen  and 
oxygen  may  be  made  clear,  it  will  be  of  assistance  to 
compare  the  oxidation  of  hydrogen  sulphide,  ammonia, 
phosphine,  and  methane.  In  the  cases  of  hydrogen  sul- 
phide and  ammonia  the  oxidation  appears  to  take  place 
as  represented  below  : 


Hydrogen  Unknown  Sulphurous  Sulphuric 

Sulphide  acid  acid 

(376) 


RELATIONS  BETWEEN  COMPOUNDS  OF  CARBON.    377 

(OH  (OH 

N^OH,    ON -{OH. 
(OH  (OH 

Yields  hypo-       Yields  nitrous        Yields  nitric 
nitrous  acid  acid  acid 


The  last  three  products,  if  formed,  break  down,  losing 
water,  and  forming  respectively  hyponitrous,  nitrous,  and 
nitric  acids.  The  change  to  hyponitrous  acid  does  not 
appear  to  be  of  an  altogether  simple  kind.  The  changes 
to  nitrous  and  nitric  acids,  however,  are  apparently  of  a 
kind  which  we  are  constantly  meeting  with,  as  has  al- 
ready been  pointed  out  (see  pp.  261-265). 

It  is  possible  that  the  changes  involved  in  the  gradual 
oxidation  of  ammonia  take  place  primarily  just  as  in  the 
oxidation  of  sulphur,  the  nitrogen  first  becoming  satu- 
rated. According  to  this  view  the  changes  should  be 
represented  as  follows  : 


H  (H 

H,  ON-{  H, 

H          H 


If  nitrous  acid  were  formed  by  the  breaking  down  of 

(OH 

the  compound  ON  •<  OH  ,  its  structure  would  probably 


be  that  represented  by  the   formula   N  •<  O  ,  or  O2NH, 

(  H 

the  hydrogen  being  in  combination  with  nitrogen  and  not 
with  oxygen.  Some  facts  seem  to  show  that  this  view  is 
probable.  While  these  changes  cannot  be  followed  very 
well  in  the  case  of  the  compounds  of  nitrogen,  and  there 
is,  therefore,  considerable  speculation  in  what  has  just 
been  said  regarding  them,  the  case  of  phosphorus  is  much 
clearer,  as  has  already  been  shown.  Here,  starting  with 
phosphine,  the  changes  are  apparently  correctly  repre- 
sented by  the  following  formulas  : 


378 


INORGANIC  CHEMISTRY. 


(H 

P^H, 
(H 

Phosphine 


(H  (OH 

OP^H,  OP^H    , 
(H  (H 


(OH  (OH 

OP^OH,  OP^OH. 
(H  (OH 


Unknown 


Hypophosphor-     Phosphorous 
ous  acid  acid 


Phosphoric 
acid 


With  methane  the  changes  effected  by  oxidation  are 
apparently  perfectly  analogous  to  those  considered.  We 
should  expect  to  find  the  following : 


OH 
OH 
OH 
OH 


But  the  tendency  of  the  compounds  containing  two 
hydroxyl  groups  to  break  down,  yielding  water  as  one  of 
the  products,  is  as  marked  as  in  the  compounds  of  nitro- 
gen. Consequently  the  products  3,  4,  and  5  do  not  exist 
in  the  free  state.  They  break  down  as  represented  in 
the  following  equations : 


(H 
CJH    +    H,0; 


=     C 


H,0; 


=    CJg    +    2H,0. 


The    products    actually  obtained,  therefore,  are  as 
follow : 


RELATIONS  BETWEEN  COMPOUNDS  OF  CARBON.    379 


(O 

|o 


Methane  Methyl  Formic  Formic  Carbon 

alcohol  aldehyde  acfd  dioxide 


By  oxidation  of  formic  acid  we  should  expect  the  for- 

(OH 
mation  of  a  product,  C  •<  O    .     While  this  is  not  known, 

(OH 

salts  of  an  acid  of  this  composition  are  known.  It  is 
ordinary  carbonic  acid,  which  when  set  free  breaks  down 
into  carbon  dioxide  and  water. 

It  will  be  seen  from  the  above  considerations  that 
there  is  a  general  analogy  between  the  changes  which 
take  place  in  passing  by  oxidation  from  the  lowest  re- 
duction-products of  the  elements  to  their  highest  oxida- 
tion-products. 

The  intermediate  stages  have  been  studied  with  special 
care  in  the  case  of  carbon  ;  the  intermediate  products  rep- 
resent classes  of  compounds  vrhich  are  not  met  with  among 
any  derivatives  of  the  other  elements.  The  intermediate 
products  in  the  case  of  sulphur  and  phosphorus  are  all 
acids.  One  of  the  intermediate  products  in  the  case  of 
nitrogen,  hydroxylamine,  is  basic,  while  the  rest  are  acid. 
The  first  oxidation-product  of  marsh-gas,  methyl  alcohol, 
CH3.OH,  is  somewhat  basic,  but  in  some  respects  differs 
from  ordinary  bases.  It  is  the  simplest  representative 
of  a  great  class  of  compounds  of  carbon  known  as  alco- 
hols, of  which  our  ordinary  alcohol,  or  spirits  of  wine,  is 
the  best  known  example.  The  next  product,  or  formic 
aldehyde,  which  has  the  structure  represented  by  the 

(H 

formula  C  1  H  or  H2C=O,  is  the  simplest  representative 

(O 

of  another  great  class  of  compounds  of  carbon  known  as 
aldehydes.  The  aldehydes  are  neither  acid  nor  basic,  but 
are  easily  converted  into  acids  by  oxidation,  and  into  bases 


380  .INORGANIC  CHEMISTRY. 

by  reduction.    The  third  oxidation-product  of  the  formula 

,H  H 

I  •"•  i 

C  <  OH   or    O=C      is  the  simplest  representative  of  a 

(°  6n 

great  class  of  carbon  compounds  known  as  the  organic 
acids.  It  is  called  formic  acid.  It  would  lead  too  far  to 
pursue  this  subject  now.  The  relations  referred  to  are 
studied  under  the  head  of  Organic  Chemistry,  or  the 
Chemistry  of  the  Compounds  of  Carbon.  Only  the  sim- 
plest oxygen  compounds  will  be  taken  up  here. 

Carbon  Dioxide,  CO2. — The  principal  compound  of  car- 
bon and  oxygen  is  carbon  dioxide,  CO2,  commonly  known 
as  carbonic  acid  gas.  Under  the  head  of  The  Air  atten- 
tion was  called  to  the  fact  that  this  gas  is  a  constant  con- 
stituent of  the  air,  though  its  relative  quantity  is  small — 
about  3  parts  in  10,000.  It  issues  from  the  earth  in  many 
places,  particularly  in  the  neighborhood  of  volcanoes. 
Many  mineral  waters  contain  it  in  large  quantity,  promi- 
nent among  which  are  the  waters  of  Pyrmont,  Selters,  and 
the  Geyser  Spring  of  Saratoga.  In  small  quantity  it  is 
present  in  all  natural  waters.  In  combination  with  bases 
it  occurs  in  enormous  quantities,  particularly  in  the  form 
of  calcium  carbonate,  CaCO3,  varieties  of  which  are 
ordinary  limestone,  chalk,  marble,  and  calc-spar.  Dolo- 
mite, which  forms  mountain-ranges,  being  particularly 
abundant  in  the  Swiss  Alps,  is  a  compound  containing 
calcium  carbonate  and  magnesium  carbonate,  MgCO3. 

Carbon  dioxide  is  constantly  formed  in  many  natural 
processes.  Thus,  all  animals. that  breathe  in  the  air  give 
off  carbon  dioxide  from  the  lungs.  That  the  gases  from 
the  lungs  contain  carbon  dioxide  can  easily  be  shown  by 
passing  them  through  lime-water,  when  a  precipitate  of 
calcium  carbonate  is  formed. 

That  carbon  dioxide  is  formed  in  the  combustion  of 
charcoal  and  wood  has  already  been  shown.  In  a  similar 
way  it  can  be  shown  that  the  gas  is  formed  whenever  any 
of  our  ordinary  combustible  substances  are  burned. 
From  our  fires,  as  from  our  lungs,  and  from  the  lungs  of 
all  animals,  then,  carbon  dioxide  is  constantly  given  off. 


CARBON  DIOXIDE.  381 

Further,  the  natural  processes  of  decay  of  both  vegetable 
and  animal  matter  tend  to  convert  the  carbon  of  this 
matter  into  carbon  dioxide,  which  then  finds  its  way 
principally  into  the  air.  The  process  of  alcoholic  fer- 
mentation, and  some  other  similar  processes,  also  give 
rise  to  the  formation  of  carbon  dioxide.  In  all  fruit- 
juices  there  is  contained  sugar.  When  the  fruits  ripen, 
fall  to  the  earth,  and  undergo  spontaneous  change,  the 
sugar  is  converted  into  alcohol  and  carbon  dioxide.  We 
see,  thus,  that  there  are  many  important  sources  of 
supply  of  carbon  dioxide,  and  it  will  be  readily  under- 
stood why  the  gas  should  be  found  everywhere  in  the  air. 
Preparation.  —  The  easiest  way  to  get  carbon  dioxide 
not  mixed  with  other  substances  is  by  adding  an  acid  to 
a  salt  of  carbonic  acid  or  a  carbonate.  In  the  decompo- 
sition of  the  carbonates  by  other  acids  we  see  exemplified 
the  same  principle  as  that  which  is  involved  in  setting 
nitric  acid  free  from  a  nitrate,  or  hydrochloric  acid  from 
sodium  chloride  by  sulphuric  acid,  and  more  particularly 
in  the  liberation  of  sulphur  dioxide  from  a  sulphite. 
In  all  these  cases  the  products  are  volatile,  and  there- 
fore, when  a  non-volatile  acid  is  added  to  the  salts, 
decomposition  takes  place.  Sulphites  do  not  yield  the 
corresponding  acid,  but  this  breaks  down  into  water^and 
the  anhydride  : 


j  2NaN03  +  H2SO4  =  Na2SO4  +  2HNO8  ; 
|2NaCl     +  HaSO4  =  NaaS04  +  2HC1. 

]  Na2S03   +  H2SO4  =  Na2SO4  +  H2SO3  ; 
|H2S03     =  S02      +H20. 

(  Na2C03   +  H0S04  =  Na2S04  +  H2CO3  ; 
|H2C03     =C02     +HaO. 

Any  acid  which  is  not  volatile  at  the  ordinary  tempera- 
ture will  decompose  a  carbonate  and  cause  an  evolution 
of  carbon  dioxide.  The  action  between  sodium  carbon- 
ate and  hydrochloric  acid  is  represented  in  this  way  : 

Na2C08  +  2HC1  =  2NaCl  +  CO3  +  HaO  ; 


382  INORGANIC  CHEMISTRY. 

that  between  nitric  acid  and  sodium  carbonate  in  this 
way : 

NaaCO3  +  2HNO3  =  2NaNO,  +  CO,  +  H2O. 

For  the  purpose  of  preparing  carbon  dioxide  in  the 
laboratory,  calcium  carbonate,  in  the  form  of  marble  or 
limestone,  and  hydrochloric  acid  are  commonly  used. 
The  reaction  involved  is  represented  thus  : 

CaCO3  +  2HC1  =  CaCl2  +  CO3  +  H2O. 

The  apparatus  used  is  the  same  as  that  used  in  making 
hydrogen  from  zinc  and  sulphuric  acid.  As  the  gas  is 
somewhat  soluble  in  water,  it  is  best  for  ordinary  pur- 
poses to  collect  it  by  displacement  of  air,  the  vessel  being 
placed  with  the  mouth  upward,  as  the  gas  is  considerably 
heavier  than  air. 

Properties. — Carbon  dioxide  is  a  colorless  gas  at  or- 
dinary temperatures.  When  subjected  to  a  low  tempera- 
ture and  high  pressure  it  is  converted  into  a  liquid. 
Liquid  carbon  dioxide  is  now  manufactured  on  the  large 
scale  for  use  as  a  fire-extinguisher,  and  for  the  purpose 
of  charging  liquids  with  the  gas.  When  some  of  the 
liquid  is  exposed  to  the  air  evaporation  takes  place  so 
rapidly  that  a  great  deal  of  heat  is  absorbed,  and  some 
of  the  liquid  becomes  solid.  The  gas  has  a  slightly  acid 
taste  and  smell.  It  is  not  combustible,  nor  does  it  sup- 
port combustion.  It  is  not  combustible  for  the  same 
reason  that  water  is  not :  because  it  already  holds  in 
combination  all  the  oxygen  it  has  the  power  to  combine 
with.  Before  it  can  burn  again  it  must  first  be  decom- 
posed. As  regards  the  statement  that  it  does  not  support 
combustion,  it  should  be  remarked  that  this  is  only  rela- 
tively true.  The  compound  does  not  easily  give  up 
oxygen,  but  to  some  substances  it  does  give  it  up,  and 
some  such  substances  burn  in  it.  For  example,  the  ele- 
ment potassium,  which,  as  we  have  seen,  has  the  power 
to  decompose  water,  has  also  the  power  to  decompose 
carbon  dioxide  if  heated  in  it  to  a  sufficiently  high  tem- 
perature, and  when  the  decomposition  once  begins,  it 


RESPIRATION.  383 

proceeds  with  brilliancy,  the  act  being  accompanied  by 
a  marked  evolution  of  heat  and  light.  Carbon  dioxide 
is  much  heavier  than  air,  its  specific  gravity  being  1.529. 
A  liter  of  the  gas  under  standard  conditions  of  tempera- 
ture and  pressure  weighs  1.977  grams.  It  dissolves  in 
water,  one  volume  of  water  dissolving  about  one  volume 
of  the  gas  at  the  ordinary  temperature.  As  is  the  case 
with  all  gases,  when  the  pressure  is  increased  the  water 
dissolves  more  gas,  and  when  the  pressure  is  removed 
the  gas  again  escapes.  The  so-called  "  soda-water"  is 
simply  water  charged  with  carbon  dioxide  under  pressure. 
The  escape  of  the  gas  when  the  water  is  drawn  is  famil- 
iar to  every  one.  The  name  soda-water  has  its  origin  in 
the  fact  that  the  carbon  dioxide  used  in  charging  the 
water  is  frequently  made  from  primary  or  acid  sodium 
carbonate,  NaHCO3,  which  is  also  called  soda  or  bicar- 
bonate of  soda. 

Relations  of  Carbon  Dioxide  to  Chemical  Energy. — 
Carbon  has  the  power  to  combine  with  oxygen,  and  in  so 
doing  a  definite  quantity  of  heat  is  evolved.  A  kilogram 
of  carbon  represents  a  certain  quantity  of  chemical 
energy,  which  we  can  get  from  it  first  in  the  form  of  heat, 
and  by  transformation,  in  other  forms  of  energy,  as  mo- 
tion, electrical  energy,  etc.  After  the  kilogram  of  carbon 
has  been  burned  it  no  longer  represents  the  energy  it 
did  in  the  form  of  carbon.  A  body  of  water  elevated 
ten  or  fifteen  feet  represents  a  certain  quantity  of  energy 
which  can  be  obtained  by  allowing  the  water  to  fall  upon 
the  paddles  of  a  water-wheel  connected  with  the  machin- 
ery of  a  mill.  After  the  water  has  fallen,  however,  it  no 
longer  has  the  power  to  do  work,  or  it  has  none  of  the 
energy  which  it  possessed  by  virtue  of  its  position.  In 
order  that  it  may  again  do  work  it  must  again  be  lifted. 
So,  too,  in  order  that  the  carbon  in  carbon  dioxide  may 
again  do  work  the  compound  must  be  decomposed. 

Respiration. — It  was  stated  above  that  carbon  dioxide 
is  given  off  from  the  lungs  just  as  it  is  from  a  fire.  It  is 
a  waste-product  of  the  processes  taking  place  in  the  ani- 
mal body.  Just  as  it  cannot  support  combustion,  so  also 
it  cannot  support  respiration.  It  is  not  poisonous  any 


384  TNOEGANIC  CHEMISTRY. 

more  than  water  is  ;  but  it  is  not  able  to  supply  the  oxy* 
gen  which  is  needed  for  breathing  purposes,  and  hence 
animals  die  when  placed  in  it.  They  die  by  suffocation, 
very  much  as  they  do  in  drowning.  Any  considerable 
increase  in  the  quantity  of  carbon  dioxide  in  the  air  above 
that  which  is  normally  present  is  objectionable,  for  the 
reason  that  it  decreases  the  proportion  of  oxygen  in  the 
air  which  is  breathed.  If,  however,  pure  carbon  dioxide 
is  introduced  into  the  air,  it  has  been  found  that  as  much 
as  5  per  cent  may  be  present  without  serious  results  to 
those  who  breathe  it.  In  a  badly  ventilated  room  in 
which  a  number  of  people  are  collected,  and  lights  are 
burning,  it  is  well  known  that  in  a  short  time  the  air  be- 
comes foul,  and  bad  effects,  such  as  headache,  drowsi- 
ness, etc.,  are  produced  on  the  occupants  of  the  room. 
These  effects  appear  to  be  due,  not  to  the  carbon 
dioxide,  but  largely  to  other  waste-products  which  are 
given  off  from  the  lungs  in  the  process  of  breathing. 
The  gases  given  off  from  the  lungs  consist  of  nitrogen, 
oxygen,  carbon  dioxide,  and  water  vapor.  Besides  these,, 
however,  there  are  many  substances  in  small  quantity, 
in  a  finely  divided  condition,  which  contain  carbon,  and 
are  in  a  state  of  decomposition.  These  act  as  poisons, 
and  they  are  the  chief  cause  of  the  bad  effects  experi- 
enced in  breathing  air  which  is  contaminated  by  the  exha- 
lations from  the  lungs.  As  carbon  dioxide  is  given  off 
from  the  lungs  at  the  same  time,  the  quantity  of  this 
gas  present  is  proportional  to  the  quantity  of  the  or- 
ganic impurities.  Hence,  by  determining  the  quantity 
of  carbon  dioxide  it  is  possible  to  form  an  opinion  a& 
to  whether  the  air  of  a  room  occupied  by  human  beings 
is  fit  for  use  or  not. 

As  carbon  dioxide  is  formed  in  the  earth  wherever  an 
acid  solution  comes  in  contact  with  a  carbonate,  the  gas 
is  frequently  given  off  from  fissures  in  the  earth.  It  is 
hence  not  unfrequently  found  in  old  wells  which  have  not 
been  in  use  for  some  time,  and  deaths  have  been  caused 
by  descending  these  wells  for  the  purpose  of  repairing 
them.  The  gas  is  also  frequently  met  with  in  mines,  and 
is  called  choke-damp  by  the  miners.  The  miners  are 


CARBON  DIOXIDE  AND  LIFE.  385 

;iware  that  after  an  explosion  caused  by  fire-damp  there 
is  danger  of  death  from  choke-damp.  The  reason  of  the 
presence  of  this  gas  after  an  explosion  is  simple.  When 
fire-damp,  or  marsh-gas,  explodes  with  air  the  carbon  is 
oxidized  to  choke-damp,  or  carbon  dioxide,  and  the  hy- 
drogen to  water.  Air  in  which  a  candle  will  not  burn 
is  not  fit  for  breathing  purposes. 

Carbon  Dioxide  and  Life. — The  role  played  by  carbon 
dioxide  in  nature  is  extremely  important  and  interesting. 
The  carbon  contained  in  living  things  is  obtained  from 
carbon  dioxide,  and  generally  returns  to  this  form  when 
life  ceases.  We  have  seen  that  all  living  things  contain 
carbon  as  an  essential  constituent.  Whence  comes  this 
carbon?  Animals  eat  either  the  products  of  plant-life 
or  other  animals  which  derive  their  sustenance  from  the 
vegetable  kingdom.  The  food  of  animals  comes,  then, 
either  directly  or  indirectly  from  plants.  But  plants 
derive  their  sustenance  largely  from  the  carbon  dioxide 
of  the  air.  The  plants  have  the  power  to  decompose  the 
gas  with  the  aid  of  the  direct  light  of  the  sun,  and  they 
then  build  up  the  complex  compounds  of  carbon  which 
form  their  tissues,  using  for  this  purpose  the  carbon  of 
the  carbon  dioxide  which  they  decompose.  Many  of 
these  compounds  are  fit  for  food  for  animals ;  that  is  to 
say,  they  are  of  such  composition  that  the  forces  at  work 
in  the  animal  body  are  capable  of  transforming  them 
into  animal  tissues,  or  of  oxidizing  them,  and  thus  keep- 
ing the  temperature  of  the  body  up  to  the  necessary 
point.  That  part  of  the  food  which  undergoes  oxidation 
in  the  body  plays  the  same  part  as  fuel  in  a  stove.  It  is 
burned  up  with  an  evolution  of  heat,  the  carbon  being 
converted  into  carbon  dioxide,  which  is  given  off  from 
the  lungs.  From  fires  and  from  living  animals  carbon 
dioxide  is  returned  to  the  air,  where  it  again  serves  as 
food  for  plants.  When  the  life  process  stops  in  the  ani- 
mal or  the  plant,  decomposition  begins ;  and  the  final 
result  of  this,  under  ordinary  circumstances,  is  the  con- 
version of  the  carbon  into  the  dioxide. 

Energy  Stored  up  in  Plants. — It  will  thus  be  seen  that 
under  the  influence  of  life  and  sunlight  carbon  dioxide  is 


386  INORGANIC  CHEMISTRY. 

constantly  being  converted  into  compounds  containing 
carbon  which  are  stored  up  in  the  plant.  These  com- 
pounds are  capable  of  burning,  and  thus  giving  heat ;  or 
some  of  them  may  be  used  as  food  by  animals,  when  they 
assume  other  forms  under  the  influence  of  the  life-process 
of  the  animals.  As  long  as  life  continues,  plants  and 
animals  are  storehouses  of  energy.  When  death  occurs, 
the  carbon  compounds  begin  to  pass  back  to  the  form  of 
carbon  dioxide,  and  the  chemical  energy  is  transformed 
partly  into  heat,  and  is  thus,  as  we  say,  dissipated.  The 
power  to  do  work,  which  the  carbon  compounds  of  plants 
and  animals  possess,  comes  from  the  heat  of  the  sun. 
It  takes  a  certain  quantity  of  this  heat,  operating  under 
proper  conditions,  to  decompose  a  certain  quantity  of 
carbon  dioxide,  and  elaborate  the  compounds  contained 
in  the  plants.  When  these  compounds  are  burned  they 
give  out  the  heat  which  was  absorbed  in  their  formation 
during  the  growth  of  the  plants.  These  compounds  are 
said  to  possess  chemical  energy.  This  has  its  origin  in 
heat,  and  is  capable  of  reconversion  into  heat.  The 
transformation  of  the  energy  of  the  sun's  heat  into  chemi- 
cal energy  lies  at  the  foundation  of  all  life.  As  the  heat 
of  the  sun  acting  upon  the  great  bodies  of  water  and  on 
the  air  gives  rise  to  the  movements  of  water  which  are 
so  essential  to  the  existence  of  the  world  as  it  is,  so  the 
action  of  the  sun's  rays  on  carbon  dioxide,  under  the 
influence  of  the  delicate  and  inexplicable  mechanism  of 
the  leaf  of  the  plant,  gives  rise  to  those  changes  in  the 
forms  of  combination  of  the  element  carbon  which  ac- 
company and  are  fundamental  to  the  wonderful  process 
of  life. 

Carbonic  Acid  and  Carbonates. — When  carbon  dioxide 
is  passed  into  water  the  solution  has  a  slightly  acid  re- 
action. The  solution  will  act  upon  bases  and  form  salts. 
The  formula  of  the  sodium  salt  formed  in  this  way  has 
been  shown  to  be  Na2CO3 ;  that  of  the  potassium  sail, 
K2CO8 ;  etc.  These  salts  are  plainly  derived  from  an 
acid,  H2CO3,  which  is  called  carbonic  acid.  It  is  prob- 
able that  this  acid  is  contained  in  the  solution  of  carbon 


CARBONIC  ACID  AND   CARBONATES.  387 

dioxide  in  water.     It  is,  however,  so  unstable  that  it 

breaks  up  into  carbon  dioxide  and  water : 

-  -~~-.. 

H2C03  =  C02  +  H20. 

The  formation  of  a  salt  by  the  action  of  carbon  di- 
oxide on  a  base  takes  place  as  shown  in  the  following 
equations  : 

2KOH    +  CO2  =  K2CO3  +  H2O  ; 
Ca(OH)2  +  COa  =  CaGO,  +  H3O. 

"With  the  acid  the  action  would  take  place  as  represented 
thus : 

2KOH    +  H2C03  =  K2CO3  +  2H2O  ; 

Ca(OH)2  +  H2C03  =  CaC03  +  2H2O. 

There  is  perfect  analogy  between  the  action  of  carbon 
dioxide  and  that  of  sulphur  dioxide  on  basic  solutions. 
With  potassium  hydroxide  and  calcium  hydroxide,  sul- 
phur dioxide  acts  as  represented  in  the  following  equa- 
tions : 

2KOH    +  SO2  =  K2SO3  +  H2O ; 
Ca(OH)2  +  SO,  =  CaS03  +  H2O. 

The  products  formed  are  sulphites  or  salts  of  sulphur- 
ous acid. 

Like  sulphurous  acid,  carbonic  acid  is  dibasic,  and 
forms  two  series  of  salts,  the  primary  and  secondary, 
or  the  acid  and  normal  salts.  The  primary  or  acid  salts 
have  the  general  formula  HMCO3,  and  the  secondary  or 
normal  salts  have  the  general  formula  M2CO3.  Exam- 
ples of  the  former  are  HKCO3,  HNaCO3,  CaH2(CO3)2, 
etc. ;  and  of  the  latter  K.CO,,  Na2CO3,  CaCO3,  BaCO3, 
etc.  It  also  readily  forms  basic  salts,  as,  for  example, 
basic  copper  carbonate.  Neutral  copper  carbonate  is  to 
be  regarded  as  formed  by  the  action  of  one  molecule  of 
the  dibasic  carbonic  acid  upon  one  molecule  of  the  di- 
acid  copper  hydroxide,  Cu(OH)2 : 


388  INORGANIC  CHEMISTRY. 

OC<OH  +  HO>Cu  =  oc<o>Cu 

One  of  the  simplest  basic  carbonates  of  copper  is  that 
formed  by  the  action  of  two  molecules  of  copper  hydrox- 
ide upon  one  molecule  of  carbonic  acid : 

or    OH   ,   (HO)Cu(OH)       op    OCuOH   ,oH  o 
00<OH  +  (HO)Cu(OH)  =  00<OOuOH  +  2H*°- 

Another  basic  salt  of  more  complicated  composition  is 
that  of  magnesium.  It  is  to  be  regarded  as  derived  from 
carbonic  acid  as  represented  in  this  formula : 


There  are  some  salts  which  are  derived  from  a  pyro- 
carbonic  acid,  that  is,  a  form  of  the  acid  derived  from 
two  molecules  of  the  acid  by  loss  of  one  molecule  of 
water : 

200  <  OH  =  H2C205  +  H20. 

Such  a  salt  is  formed  by  loss  of  water  from  the  primary 
sodium  salt : 

2HNaCOs  =  Na2C205  +  H20. 

There  are  no  salts  known  derived  from  normal  carbonic 
acid,  C(OH)4,  though  there  are  some  compounds  analo- 
gous to  salts  which  are  derivatives  of  this  normal  acid. 
The  secondary  or  normal  salts  which  carbonic  acid  forms 
with  the  most  strongly  marked  metallic  elements,  viz., 
potassium  and  sodium,  are  not  decomposed  by  heat, 
but  all  other  carbonates  are  decomposed  by  heat  more 
or  less  easily,  according  to  the  strength  of  the  base. 
Calcium  carbonate  when  ignited  loses  carbon  dioxide, 
and  lime,  or  calcium  oxide,  remains  behind : 

CaCO3  =  CaO  +  CO2. 


CARBON  MONOXIDE.  389 

Carbon  Monoxide,  CO. — When  a  substance  containing 
carbon  burns  in  an  insufficient  supply  of  air, — as,  for 
example,  when  the  draught  in  a  furnace  is  not  strong 
enough  to  remove  the  products  of  combustion  and  sup- 
ply fresh  air, — the  oxidation  of  the  carbon  is  not  com- 
plete, and  the  product,  instead  of  being  carbon  dioxide, 
is  carbon  monoxide,  CO.  This  compound  can  also  be 
made  by  extracting  oxygen  from  carbon  dioxide.  It  is 
only  necessary  to  pass  the  dioxide  over  heated  carbon, 
when  reaction  takes  place  as  represented  thus  : 

CO2  +  C  =  2CO. 

This  method  of  formation  is  illustrated  in  coal  fires,  and 
can  be  well  observed  in  an  open  grate.  The  air  has  free 
access  to  the  coal,  and  at  the  surface  complete  oxidation 
takes  place.  But  that  part  of  the  carbon  dioxide  which 
is  formed  at  the  lower  part  of  the  grate  is  drawn  up 
through  the  heated  coal,  and  is  partly  reduced  to  carbon 
monoxide.  When  the  monoxide  escapes  from  the  upper 
part  of  the  grate  it  again  combines  with  oxygen,  or  burns, 
giving  rise  to  the  characteristic  blue  flame  always  noticed 
above  a  mass  of  burning  anthracite  coal.  Should  any- 
thing occur  to  prevent  free  access  of  air,  carbon  monox- 
ide may  readily  escape  complete  oxidation. 

The  monoxide  is  also  formed  by  passing  steam  over 
highly  heated  carbon,  when  this  reaction  takes  place : 

C  +  H20  =  CO  +  H3. 

This  is  the  reaction  made  use  of  in  the  manufacture  of 
"  water-gas."  The  gas  thus  obtained  is  largely  a  mixture 
of  hydrogen  and  carbon  monoxide.  The  gas  is  enriched 
by  passing  it  through  a  furnace  in  which  it  is  mixed  with 
highly  heated  vapors  of  hydrocarbons  from  petroleum. 
The  main  reaction,  the  decomposition  of  water  by  heated 
carbon,  is  effected  in  large  furnaces  filled  with  anthracite 
coal.  The  coal  is  first  heated  to  a  high  temperature 
by  setting  fire  to  it,  the  products  of  combustion  being 
allowed  to  escape.  When  it  is  hot  enough,  the  air  is 


390  INORGANIC  CHEMISTRY. 

shut  off  and  steam  passed  rapidly  in,  when  the  decom- 
position of  the  water  by  the  carbon  takes  place.  Soon 
the  mass  becomes  so  much  cooled  that  the  reaction 
stops.  The  steam  is  then  cut  off  and  air  turned  on  again, 
and  so  on. 

The  easiest  way  to  make  carbon  monoxide  is  by  heat- 
ing oxalic  acid,  which  is  a  compound  of  carbon,  hydro- 
gen, and  oxygen,  of  the  formula  C2H2O4,  with  five  to  six 
times  its  weight  of  concentrated  sulphuric  acid.  The 
change  which  takes  place  is  represented  thus : 

C2H304  =  C02  +  CO  +  H20. 

Both  the  dioxide  and  monoxide  of  carbon  are  formed. 
Both  are  gases.  In  order  to  separate  them  the  mixture 
is  passed  through  a  solution  of  sodium  hydroxide,  which 
takes  up  the  carbon  dioxide,  forming  sodium  carbonate, 
and  allows  the  monoxide  to  pass. 

Carbon  monoxide  is  a  colorless,  tasteless,  inodorous 
gas,  insoluble  in  water.  It  burns  with  a  pale-blue  flame, 
forming  carbon  dioxide.  It  is  exceedingly  poisonous 
when  inhaled.  Hence  it  is  very  important  that  it  should 
not  be  allowed  to  escape  into  rooms  occupied  by  human 
beings.  We  not  unfrequently  hear  of  deaths  caused  by 
the  gases  from  coal  stoves.  The  most  dangerous  of  the 
gases  given  off  from  these  stoves  is  probably  carbon 
monoxide.  A  pan  of  smouldering  charcoal  gives  off  this 
gas,  and  the  fact  that  it  is  poisonous  is  well  known.  It 
has  been  used  to  a  considerable  extent  for  the  purpose 
of  suicide.  The  poisonous  character  of  carbon  monoxide 
has  led  to  a  great  deal  of  discussion  and  to  some  legisla- 
tion on  the  subject  of  "  water-gas."  The  question  has 
been  repeatedly  raised  whether  government  should  allow 
the  manufacture  of  the  gas.  There  is  no  doubt  of  the 
fact  that  it  is  a  dangerous  substance,  and  that  it  should 
not  be  allowed  to  escape  into  the  air  is  obvious.  Where- 
ever  it  is  used  special  precautions  should  be  taken  to 
guard  against  leaking.  There  is  no  doubt  that  it  is 
somewhat  more  poisonous  than  coal-gas. 

At  high  temperatures  carbon  monoxide  has  a  very 


FORMIC  ACID.  391 

strong  tendency  to  combine  with  oxygen,  and  is  hence  a 
good  reducing  agent.  In  the  reduction  of  iron  from  its 
ores,  the  carbon  monoxide  formed  in  the  blast-furnace 
plays  an  important  part  in  the  reducing  process.  At 
ordinary  temperatures  the  gas  does  not  combine  readily 
with  oxygen.  Thus,  it  does  not  act  upon  ozone,  even 
when  heated  with  it  somewhat  above  the  temperature  at 
which  ozone  is  converted  into  ordinary  oxygen.  When 
passed  over  some  substances  which  are  rich  in  oxygen, 
as,  for  example,  chromic  anhydride,  CrO3,  and  potassium 
permanganate,  KMnO4,  in  acid  solution,  it  takes  up  oxy- 
gen even  at  the  ordinary  temperature.  It  unites  with 
chlorine  in  the  direct  sunlight,  and  forms  the  com- 
pound known  as  carbonyl  chloride,  COC12. 

Formic  Acid,  H2CO2. — Just  as  carbon  dioxide  may  be 
regarded  as  the  anhydride  of  carbonic  acid,  so  carbon 
monoxide  may  be  regarded  as  the  anhydride  of  an  acid 
of  the  formula  H2CO2.  While,  however,  an  acid  of  this 
formula  is  known,  it  is  not  formed  by  action  of  carbon 
monoxide  upon  water,  nor  are  its  salts  easily  formed  by 
the  action  of  carbon  monoxide  upon  bases.  By  passing 
it  over  certain  basic  substances,  however,  as,  for  example, 
potassium  hydroxide  and  calcium  hydroxide,  at  a  com- 
paratively high  temperature  action  takes  place,  and  salts 
of  the  acid  are  formed.  Thus,  in  the  case  of  potas- 
sium hydroxide,  the  action  takes  place  as  represented  in 
the  equation 

CO  +  KOH  =  HCO2K. 

Although  it  contains  two  atoms  of  hydrogen  in  the 
molecule,  formic  acid  is  monobasic.  This  fact  finds  its 
explanation  in  the  structure  of  the  acid.  All  its  reactions 
show  that  only  one  of  the  hydrogen  atoms  of  formic 
acid  is  in  combination  with  oxygen,  while  the  other  is  in 
combination  with  carbon,  as  represented  in  the  formula 

?  (H 

HC-OH  or  C  K  O     .     The  relations  here  are  similar  to 

(OH 
those  met  with  in  phosphorous  and  sulphurous  acids, 


392  i        INORGANIC  CHEMISTRY. 

which  have  been  so  frequently  referred  to.  Formic 
acid  bears  to  carbonic  acid  the  same  relation  that 
sulphurous  bears  to  sulphuric  acid,  and  phosphorous 
to  phosphoric  acid,  as  shown  in  the  formulas  : 


Formic  acid  Carbonic  acid 


o,sj 


H  0  q ( OH 

OH  u'b  \  OH 

Sulphurous  acid  Sulphuric  acid 


(H  (OH 

OP  \  OH  OP  \  OH 

(OH  (OH 

Phosphorous  acid  Phosphoric  acid 

Formic  acid  occurs  in  nature  in  red  ants,  in  stinging 
nettles,  and  elsewhere.  It  is  a  colorless  liquid,  which 
solidifies  at  8°.  6.  When  treated  with  concentrated  sul- 
phuric acid  it  breaks  down  into  carbon  monoxide  and 
water : 

H3CO,  =  CO  +  HaO. 

By  oxidation  it  is  converted  into  carbon  dioxide  and 
water. 

Carbonyl  Chloride,  Phosgene,  COC12.— This  compound 
was  referred  to  above  as  being  formed  when  chlorine  acts 
upon  carbon  monoxide  under  the  influence  of  the  sun's 
rays.  It  is  also  formed  when  the  two  gases  are  passed 
together  through  a  tube  filled  with  pieces  of  animal  char- 
coal. It  is  a  colorless  gas,  which  is  easily  condensed  to 
a  liquid  boiling  at  8°. 2.  It  is  now  manufactured  on  the 
large  scale  for  use  in  the  preparation  of  certain  classes 
of  dye-stuffs. 

Like  the  oxychlorides  of  sulphur  and  of  phosphorus, 
this  compound,  which  is  an  oxychloride  of  carbon,  is  de- 
composed by  water,  forming  carbonic  acid  or  its  products 
of  decomposition : 


CARBONTL  CHLORIDE  OR  PHOSGENE.  393 

+  181 


Cl        HOH  (  OH 

Cl  +  HOH  =  PO^  OH  +  3HC1. 

Cl        HOH  (  OH 


It  is  interesting  to  note  that,  while  the  chlorides  of 
sulphur  and  phosphorus,  SC12  and  PC13,  as  well  as 
SC14  and  PC16,  are  easily  decomposed  by  water,  the 
tetrachloride  of  carbon,  CC14,  is  not.  On  the  other 
hand,  the  tetrachloride  is  not  formed  when  the  oxides 
of  carbon  are  treated  with  hydrochloric  acid.  It 
will  be  remembered  that,  in  discussing  the  relations  be- 
tween the  acid-forming  and  the  base-forming  elements, 
attention  was  called  to  the  fact  that,  in  general,  the  chlo- 
rides of  the  acid-forming  elements  are  easily  decom- 
posed by  water,  forming  the  corresponding  hydroxides  or 
oxides,  while  the  oxides  and  hydroxides  of  the  base-form- 
ing elements  are  converted  into  chlorides  by  treatment 
with  hydrochloric  acid.  In  carbon  we  have  an  example 
of  an  element  which  occupies  what  may  be  called  almost 
a  neutral  ground  between  the  two  classes  of  elements. 
It  forms  both  acids  and  bases,  to  be  sure,  but  these  are, 
generally  speaking,  not  as  strongly  marked  as  the  acids 
and  bases  formed  by  other  elements.  This  neutral  char- 
acter of  the  element  is  also  well  shown  in  the  conduct  of 
its  chloride  towards  water,  and  of  its  oxides  towards 
hydrochloric  acid. 


CHAPTER  XXI. 

ILLUMINATiON-FLAME— BLOW-PIPE. 
COMPOUNDS  OF  CARBON  WITH  NITROGEN  AND  SULPHUR 

Introduction. — As  the  substances  used  for  illumina- 
tion contain  carbon,  and  the  chemical  processes  involved 
consist  largely  in  the  oxidation  of  the  carbon  of  these 
compounds,  this  is  an  appropriate  place  to  consider 
briefly  the  subject  of  illumination  from  a  chemical  point 
of  view,  as  well  as  that  of  flame,  and  the  blow -pipe, 
which  gives  an  extremely  useful  form  of  flame  constantly 
used  in  the  laboratory. 

In  all  ordinary  kinds  of  illumination  we  are  dependent 
upon  flames  for  the  light.  Whether  we  use  illuminating 
gas,  a  lamp,  or  a  candle,  the  light  comes  from  a  flame. 
In  the  first  case,  the  gas  is  burned  directly  ;  in  the  case 
of  the  lamp,  the  oil  is  first  drawn  up  the  wick,  then  con- 
verted into  a  gas,  and  this  burns  ;  while,  finally,  in  the 
case  of  the  candle,  the  solid  material  of  the  candle  is 
first  melted,  then  drawn  up  the  wick,  converted  into  gas, 
and  the  gas  burns,  forming  the  flame.  In  each  case  we 
have,  then,  to  deal  with  a  burning  gas,  and  this  burning 
gas  is  called  a  flame. 

Illuminating  Gas,  Coal-gas. — Illuminating  gas  is  some- 
times made  from  coal  by  heating  in  closed  retorts.  As 
has  already  been  explained,  coal,  particularly  the  softer 
kinds,  contains  compounds  of  carbon  and  hydrogen, 
together  with  some  nitrogen  and  other  elements.  When 
it  is  subjected  to  destructive  distillation,  as  in  the  manu- 
facture of  coal-gas,  the  hydrogen  passes  off  partly  in 
combination  with  carbon,  as  hydrocarbons,  and  partly  in 
the  free  state.  The  nitrogen  passes  off  as  ammonia,  and 
a  large  percentage  of  the  carbon  remains  behind  in  the 
retort  in  the  uncombined  state  as  coke.  The  gases  given 

(394) 


FLAMES.  395 

.'  "f 

off  are  purified,  and  form  illuminating  gas.  One  ton  of 
coal  yields  on  an  average  10,000  cubic  feet  of  gas.  The 
value  of  a  gas  depends  upon  the  quantity  of  light  given 
by  the  burning  of  a  definite  quantity.  It  is  measured 
by  comparing  it  with  the  light  given  by  a  candle  burn- 
ing at  a  certain  rate.  The  standard  candle  is  one  made 
of  spermaceti,  which  burns  at  the  rate  of  120  grains  per 
hour ;  that  is  to  say,  a  candle  which,  burning  under 
ordinary  conditions,  loses  120  grains  in  one  hour.  The 
standard  burner  used  for  the  gas  is  one  through  which 
five  cubic  feet  of  gas  pass  per  hour.  Now,  to  determine 
the  illuminating  power  of  a  gas,  it  is  passed  through  the 
standard  burner  at  the  rate  mentioned,  and  the  light 
which  it  gives  is  compared  with  the  light  given  by  the 
standard  candle.  This  comparison  is  easily  made  by 
means  of  an  instrument  called  the  photometer.  The 
illuminating  power  of  the  gas  is  then  stated  in  terms  of 
the  standard  candle.  The  statement  that  the  illumi- 
nating power  of  a  gas  is  fourteen  candles,  signifies  that, 
when  burning  at  the  rate  of  five  cubic  feet  per  hour,  its 
flame  gives  fourteen  times  as  much  light  as  that  of  the 
standard  candle. 

Flames. — Ordinarily  when  we  speak  of  a  flame  wo 
mean  a  gas  which  is  combining  with  oxygen.  The  hy- 
drogen flame  is  simply  the  phenomenon  accompanying 
the  act  of  combination  of  the  two  gases  hydrogen  and 
oxygen.  Owing  to  the  fact  that  we  are  surrounded  by 
oxygen,  we  speak  of  hydrogen  as  the  burning  gas.  How 
would  it  be  if  we  were  surrounded  by  an  atmosphere  of 
hydrogen?  Plainly,  oxygen  would  then  be  a  burning 
gas.  If  we  allow  a  jet  of  oxygen  to  escape  into  a  vessel 
containing  hydrogen,  a  flame  will  appear  where  the  oxy- 
gen escapes  from  the  jet,  if  a  light  is  applied.  This  is 
an  experiment  which  requires  special  precautions,  and, 
as  the  principle  can  be  illustrated  as  well  by  means  of 
illuminating  gas,  this  may  be  used  instead.  Just  as 
illuminating  gas  burns  in  an  atmosphere  of  oxygen,  so 
oxygen  burns  in  an  atmosphere  of  illuminating  gas. 

Kindling  Temperature  of  Gases. — In  studying  the  action 
of  oxygen  upon  other  substances,  we  learned  that  it  is 


396  INORGANIC  CHEMISTRY, 

necessary  that  each  of  these  substances  should  be  raised 
to  a  certain  temperature  before  it  will  combine  with 
oxygen.  This  statement  is  as  true  of  gases  as  of  other 
substances.  When  a  current  of  hydrogen  is  allowed  to 
escape  into  the  air,  or  into  oxygen,  no  action  takes  place 
unless  it  is  heated  up  to  its  burning  temperature,  when 
it  takes  fire  and  continues  to  burn,  as  the  burning  of 
one  part  of  the  gas  heats  up  the  part  which  follows  it, 
and  hence  it  is  heated  up  to  the  burning  tempera- 
ture as  fast  as  it  escapes  into  the  air.  If  the  gas  is 
cooled  down  even  very  slightly  below  this  temperature, 
it  is  extinguished.  This  can  easily  be  shown  by  bring- 
ing down  upon  the  flame  of  a  Bunsen  burner  a  piece  of 
wire  gauze.  There  will  be  no  flame  above  the  gauze, 
but  gas  will  pass  through  unburned,  and  this  will  burn 
if  it  is  lighted  above  the  gauze.  In  this  case,  by  simply 
passing  through  the  thin  wire  gauze,  the  gases  are  cooled 
down  below  their  burning  temperatures,  and  the  flame 
does  not  pass  through.  So,  also,  if  the  gas  is  turned 
on  and  not  lighted,  and  the  gauze  held  an  inch  or  two 
above  the  outlet,  the  gas  will  burn  above  the  gauze  if 
lighted  above,  and  will  not  pass  downward  through  the 
gauze,  unless  this  becomes  very  hot. 

Miner's  Safety -lamp. — The  principle  illustrated  in  the 
experiments  referred  to  in  the  last  para- 
graph is  utilized  in  the  miner's  safety- 
lamp,  to  which  reference  has  already  been 
made.  One  of  the  dangers  which  the  coal- 
miner  has  to  encounter  is  the  occurrence 
in  the  mines  of  fire-damp,  or  methane, 
CH4,  which  with  air  forms  an  explosive 
mixture.  The  explosion  can  only  be 
brought  about  by  contact  of  flame  with 
the  mixture.  In  order  to  avoid  the  con- 
tact, the  flame  of  the  safety-lamp  is  sur- 
rounded by  wire  gauze,  as  shown  in  Fig. 
11.  When  a  lamp  of  this  kind  is  brought 
into  an  explosive  mixture  of  marsh-gas 
FIG.  11.  and  air,  the  mixture  passes  through  the 
wire  gauze  and  comes  in  contact  with  the  flame,  and  a 


STRUCTURE  OF  FLAMES.  39? 

small  explosion  or  a  series  of  small  explosions  inside 
the  gauze  occurs,  but  the  flame  of  the  burning  gas 
inside  the  wire  gauze  cannot  pass  through  and  raise  the 
temperature  of  the  gas  outside  to  the  burning  tempera- 
ture. Hence  no  serious  ,  explosion  can  take  place.  The 
flickering  of  the  flame  of  the  lamp,  and  the  occurrence 
of  small  explosions  inside,  furnish  the  miner  with  the 
information  that  he  is  in  a  dangerous  atmosphere.  While 
the  safety-lamp  does  undoubtedly  afford  much  protection, 
still  explosions  occur.  These  have  been  shown  to  be  caused 
by  the  presence  of  coal-dust  in  the  mines,  and  by  the  com- 
motion of  the  air  produced  in  blasting.  By  the  aid  of  the 
coal-dust,  and  by  sudden  and  violent  movements  of  the  air, 
it  is  possible  for  a  flame  surrounded  by  wire  gauze  to 
explode  a  mixture  of  marsh-gas  and  air  on  the  other  side 
of  the  gauze. 

Structure  of  Flames. — The  hydrogen  flame  consists  of 
.a  thin  envelope  of  burning  hydrogen  enclosing  unburned 
.gas,  and  surrounded  by  water  vapor,  which  is  the  prod- 
uct of  the  combustion.  The  structure  of  other  flames 
depends  upon  the  complexity  of  the  gases  burned,  and 
the  conditions  under  which  the  burning  takes  place.  In 
.general,  a  flame  consists  of  an  outer  envelope  of  gas 
•combining  with  oxygen,  and  hence  hot,  and  an  inner 
part  which  contains  unburned  gas,  which  is  compara- 
tively cool.  A  part  of  the  unburned  gas  is,  however, 
quite  hot,  and  it  would  combine  with  oxygen  were  it  not 
for  the  fact  that  it  is  surrounded  by  an  envelope  which 
prevents  access  of  air.  The  outer  hot  part  of  the  flame 
is  called  the  oxidizing  flame,  because  it  presents  condi- 
tions favorable  to  the  oxidation  of  substances  introduced 
into  it.  The  inner  hot  part  is  called  the  reducing  flame, 
because  it  consists  of  highly  heated  substances  which 
have  the  power  to  combine  with  oxygen ;  and  hence  many 
compounds  containing  oxygen  lose  it,  or  are  reduced, 
when  introduced  into  this  part  of  the  flame.  The  hot- 
test part  of  the  flame  is  about  half-way  between  the 
bottom  and  the  top.  Here  oxidation  is  taking  place 
most  energetically.  The  hottest  part  of  the  unburned 
gases  is  at  the  tip  of  the  dark  central  part  of  the  flame. 


398  INORGANIC  CHEMISTRY. 

In  the  flame  of  a  Bunsen  burner  the  two  parts  can  be 
easily  distinguished.  The  dark  central  part  of  the  flame 
extends  for  some  distance  above  the  outlet  of  the  burner. 
If  the  holes  at  the  base  of  the  burner  are  partly  closed, 
the  tip  of  the  central  part  of  the  flame  becomes  lumi- 
nous. This  luminous  tip  is  most  efficient  for  the  pur- 
pose of  reduction.  The  principal  parts  of 
the  flame  are  those  marked  in  Fig.  12.  The 
part  marked  b  is  the  central  cone  of  un- 
burned  gases ;  that  marked  c  is  the  lumi- 
nous tip,  the  best  part  of  the  flame  for  re- 
duction. A  is  the  envelope  of  burning 
gas.  The  hottest  part  of  the  flame  is  at  a  ; 
that  which  is  most  efficient  in  causing  oxi- 
dation is  at  d.  This  is  further  surrounded 
by  a  non-luminous  envelope  consisting  of 
the  products  of  combustion,  carbon  diox- 
ide and  water  vapor.  Certain  metals  placed 
in  the  upper  end  of  the  flame  take  up 
oxygen,  because  they  are  highly  heated  in  the  presence 
of  oxygen.  Certain  oxides  lose  their  oxygen  when  placed 
in  the  tip  of  the  central  cone,  because  the  gases  are  here 
heated  to  the  temperature  at  which  they  have  the 
power  to  combine  with  oxygen. 

Blow-pipe. — The  oxidizing  and  reducing  flames  are 
frequently  utilized  in  the  laboratory.  For  the  purpose 
of  increasing  their  efficiency  a  blow-pipe  is  used.  This  is 
simply  a  tube  with  a  convenient  mouth-piece  and  a  nozzle 
with  a  small  opening  through  which  air  is  blown  into  a 
flame  by  means  of  the  mouth.  The  blow-pipe  may  be 
used  with  the  flame  of  a  candle,  an  alcohol-lamp,  or  a 
gas-lamp.  It  is  commonly  used  with  a  gas-lamp.  By 
regulating  the  current  of  air  and  slightly  changing  the 
position  of  the  tip  of  the  blow-pipe  a  good  oxidizing 
flame  or  a  good  reducing  flame  can  be  produced.  Some 
oxides  are  very  easily  reduced  when  heated  in  the  re- 
ducing blow-pipe  flame.  Others  are  not.  We  can  fre- 
quently judge  of  the  composition  of  a  substance  by  heat- 
ing in  the  blow-pipe  flame,  and  noticing  its  conduct. 
Some  metals  are  easily  oxidized  in  the  oxidizing  flame. 


LUMINOSITY  OF  FLAMES.  399 

Some  form  characteristic  films,  or  thin  layers  of  oxides, 
on  the  substance  upon  which  they  are  heated,  which  is 
usually  charcoal ;  and,  in  some  cases,  it  is  possible  to 
detect  the  presence  of  certain  substances  by  the  color  of 
the  film  of  oxide.  The  blow-pipe  is  therefore  of  much 
value  as  affording  a  method  for  the  detection  of  the 
presence  of  certain  elements  in  mixtures  or  compounds 
of  unknown  composition.  The  chemical  principles  in- 
volved in  its  use  will  be  clear  from  what  has  already 
been  said. 

Causes  of  the  Luminosity  of  Flames. — It  is  evident  from 
what  we  have  seen  that  flames  differ  greatly  in  their 
light-giving  power.  The  hydrogen  flame,  for  example, 
though  extremely  hot,  gives  practically  no  light.  This  is 
also  the  case  with  the  flame  of  the  Bunsen  burner ; 
while,  on  the  other  hand,  the  flame  of  coal-gas,  burning 
under  ordinary  circumstances,  and  that  of  a  candle,  etc., 
give  light.  To  what  is  the  difference  due  ?  This  subject 
has  been  studied  very  thoroughly,  and  it  has  been  found 
that  there  are  several  causes  which  operate  to  make  a 
flame  give  light,  and  vice  versa.  In  the  first  place,  if  a 
solid  substance  which  does  not  burn  is  introduced  into 
a  non-luminous  flame,  a  part  of  the  heat  appears  as 
light.  This  is  seen  when  a  spiral  of  platinum  wire  is 
introduced  into  a  hydrogen  flame.  It  is  also  seen  when 
a  piece  of  lime  is  introduced  into  the  hot  non-luminous 
flame  of  the  oxyhydrogen  blow-pipe.  A  similar  cause 
operates  in  ordinary  gas-flames  to  make  them  luminous. 
There  are  always  present  particles  of  unburned  carbon, 
as  can  be  shown  by  putting  a  piece  of  porcelain  or  any 
solid  substance  into  the  flame,  when  there  will  be  de- 
posited on  it  a  layer  of  soot,  which  consists  mainly  of 
finely  divided  carbon.  In  the  flame  such  particles  are 
heated  up  to  incandescence,  or  to  the  temperature  at 
which  they  give  light.  Again,  it  has  been  found  that  a 
candle  gives  more  light  at  the  level  of  the  sea  than  it  does 
when  at  the  top  of  a  high  mountain,  as  Mount  Blanc,  on 
which  the  experiment  was  actually  performed.  This  is 
partly  due  to  a  difference  in  the  density  of  the  gases. 
Naturally,  the  denser  the  gas  the  more  active  the  com- 


400  INORGANIC  CHEMISTRY. 

bustion,  the  greater  the  heat,  and  the  brighter  the  light. 
This  last  statement  ceases  to  be  true  when  the  oxidation 
becomes  sufficient  to  burn  up  all  the  solid  particles  in 
the  flame.  If  gases,  which  in  burning  give  light,  are 
cooled  down  before  they  are  burned,  the  luminosity  is 
diminished,  and,  conversely,  non-luminous  flames  may 
be  rendered  luminous  by  heating  the  gases  before  burn- 
ing them.  Gases  which  otherwise  give  luminous  flames 
give  non- luminous  flames  when  diluted  to  a  sufficient 
extent  with  neutral  gases,  such  as  nitrogen  and  carbon 
dioxide,  which  neither  burn  nor  support  combustion. 

Bunsen  Burner. — All  the  statements  made  in  regard  to 
the  causes  of  the  luminosity  of  flames  are  based  upon 
carefully  performed  experiments.  These  experiments, 
however,  cannot,  for  the  most  part,  be  readily  repeated 
by  the  student  in  the  laboratory  in  a  satisfactory  way. 
One  constant  reminder  of  the  possibility  of  rendering  a 
luminous  flame  non-luminous,  and  vice  versa,  is  fur- 
nished by  the  burner  universally  used  in  chemical  labora- 
tories, and  called,  after  the  name  of  the  inventor,  the 
Bunsen  burner.  The  construction  of  this  burner  is  easily 
understood.  It  consists  of  a  base  and  an  upper  tube. 
The  base  is  connected  by  means  of  a  rubber  tube  with 
the  gas  supply.  The  gas  escapes  from  a  small  opening 
in  the  base,  and  passes  upward  through  the  tube.  At 
the  lower  part  of  the  tube  there  are  two  holes,  which 
may  be  opened  or  closed  by  turning  a  ring  with  two  cor- 
responding holes  in  it.  When  the  gas  is  turned  on,  it  is 
lighted  at  the  top  of  the  tube.  Air  is  at  the  same  time 
drawn  through  the  holes  at  the  base.  The  result  is  that 
the  flame  is  practically  non-luminous.  If  the  ring  at  the 
base  is  turned  so  that  the  air-holes  are  closed,  the 
flame  becomes  luminous.  The  advantage  of  the  non- 
luminous  flame  for  laboratory  use  consists  in  the  fact 
that  it  does  not  deposit  soot,  and,  at  the  same  time,  it 
is  hot. 

The  non-luminosity  of  the  flame  of  the  Bunsen  burner 
appears  to  be  due  to  several  causes :  (1)  Dilution  of  the 
gases  by  means  of  the  nitrogen  of  the  air ;  (2)  Cooling 
of  the  gases  by  the  entrance  of  the  air ;  (3)  Burning  of 


CYANOGEN.  401 

the  solid  particles  by  tlie  aid  of  the  oxygen  of  the  air 
admitted  to  the  interior  of  the  flame. 


COMPOUNDS  or  CARBON  WITH  NITROGEN  AND  WITH 
SULPHUR. 

Cyanogen,  C2N2.  —  Carbon  does  not  combine  with  ni- 
trogen under  ordinary  circumstances.  If,  however,  they 
are  brought  together  at  very  high  temperatures  in  the 
presence  of  metals,  they  combine  to  form  compounds 
known  as  cyanides.  Thus,  when  nitrogen  is  passed  over 
a  highly  heated  mixture  of  carbon  and  potassium  car- 
bonate, potassium  cyanide,  KCN,  is  formed.  Carbon 
containing  nitrogen,  as  animal  charcoal,  when  ignited 
with  potassium  carbonate,  reduces  the  carbonate,  form- 
ing potassium,  in  presence  of  which  carbon  and  nitro- 
gen combine,  forming  potassium  cyanide.  When  refuse 
animal  substances,  such  as  blood,  horns,  claws,  hair, 
etc.,  are  heated  together  with  potassium  carbonate  and 
iron,  a  substance  known  as  potassium  ferro-cyanide,  or 
yellow  prussiate  of  potash,  K4Fe(CN)6  -J-  3H2O,  is  formed. 
When  this  is  simply  heated  it  is  decomposed,  yielding 
potassium  cyanide  : 

K4Fe(CN)6  =  4KCN  +  FeC2  +  N2. 

It  is  an  easy  matter  to  make  the  mercury  salt,  Hg(CN)a, 
from  the  potassium  salt.  By  heating  mercuric  cyanide 
it  breaks  up,  yielding  metallic  mercury  and  cyanogen 
gas: 


just  as  mercuric  oxide  yields  mercury  and  oxygen  when 
heated  : 

HgO  =  Hg  +  0. 

Cyanogen  (from  KIMXVOS,  Hue)  owes  its  name  to  the  fact 
that  several  of  its  compounds  have  a  blue  color.  It  is  a 
colorless  gas,  which  is  easily  soluble  in  water  and  alco- 


402  INORGANIC  CHEMISTRY. 

hoi,  and  is  extremely  poisonous.  It  burns  with  a  purple- 
colored  flame.  In  aqueous  solution  cyanogen  soon  un- 
dergoes change,  and  a  brown  amorphous  substance  is 
deposited.  In  the  solution  are  found  hydrocyanic  acid, 
oxalic  acid,  ammonia,  and  carbon  dioxide.  The  princi- 
pal cause  of  this  decomposition  is  apparently  the  ten- 
dency of  the  nitrogen  to  combine  with  hydrogen  to  form 
the  stable  compound  ammonia,  and  of  carbon  to  com- 
bine with  hydrogen  and  oxygen  to  form  stable  com- 
pounds like  oxalic  acid  and  carbon  dioxide.  One  of  the 
chief  decompositions  which  cyanogen  undergoes  with 
water  is  that  represented  in  the  equation 


H 

The  compound  I     2     or  H2C2O4  is  oxalic  acid.     This 

kind  of  decomposition  with  water  is  characteristic  of  cy- 
anogen compounds.  It  consists,  as  will  be  seen,  in  the 
union  of  the  nitrogen  with  hydrogen  to  form  ammonia, 
and  the  union  of  the  carbon  with  oxygen  and  hydroxyl. 

Hydrocyanic  Acid,  Prussia  Acid,  tHCN.  —  This  acid, 
which  is  commonly  known  by  the  name  prussic  acid,  oc- 
curs in  nature  in  amygdalin,  in  combination  with  other 
substances,  in  bitter  almonds,  the  leaves  of  the  cherry, 
laurel,  etc.  It  is  prepared  by  decomposing  metallic  cy- 
anides with  hydrochloric  acid.  It  is  volatile  and  passes 
over.  The  action  is  represented  thus  : 

KCN  +  HC1  =  KC1  +  HON. 

It  can  also  be  made  by  treating  chloroform  with 
ammonia  : 

CHC18  +  NH3  =  HCN  +  3HC1. 

Of  course,  the  hydrochloric  acid  and  the  hydrocyanic 
acid  formed  combine  with  ammonia,  so  that  the  complete 
action  is  represented  by  this  equation  : 

CHC13  +  5NH3  =  NH4CN  +  3NH4C1. 
The  product  NH4CN  is  ammonium  cyanide. 


HYDROCYANIC  ACID.  403 

Hydrocyanic  acid  is  a  volatile  liquid,  boiling  at  26°.  1, 
and  solidifying  at  — 14°.  It  has  a  very  characteristic 
odor  suggestive  of  bitter  almonds.  It  is  extremely  poi- 
sonous. It  dissolves  in  water  in  all  proportions,  and  it  is 
such  a  solution  which  is  known  as  prussic  acid.  Pure 
hydrocyanic  acid  is  very  unstable.  By  standing,  a  brown 
substance  is  deposited  from  its  solution.  By  boiling 
with  alkalies  or  acids  it  is  converted  into  formic  acid  and 
ammonia.  This  is  another  example  of  the  tendency  of 
cyanogen  compounds  to  decompose  in  the  presence  of 
water,  yielding  ammonia  and  oxygen  compounds  of  car- 
bon. The  decomposition  of  hydrocyanic  acid  takes  place 
as  represented  in  the  equation 


HHO 


The  relations  between  chloroform,  formic  acid,  and 
hydrocyanic  acid  are  instructive.  By  replacing  all  the 
chlorine  atoms  of  chloroform  by  hydroxyl  a  compound  of 

roH 

OTT 

the  formula  C  -j  QJJ  would  be  formed  ;  but  this  would 


break  down  by  loss  of  water,  yielding  formic  acid,  C 

By  replacing  the  three  chlorine  atoms  by  one  nitrogen 
atom  hydrocyanic  acid  is  formed  ;  and  this  in  turn  when 
decomposed  in  presence  of  water  yields  formic  acid. 
These  relations  will  be  made  clear  by  the  aid  of  the  fol- 
lowing formulas  and  equations  : 

oH 


fHOH=Cj^  +  3H01; 
OH 


TTr\TT 

H  H 


404  INORGANIC  CHEMISTRY. 


Cyanic  Acid,  HCNO.  —  By  gentle  oxidation  of  a  cyan- 
ide it  is  converted  into  a  cyanate.  Thus,  by  melting 
together  potassium  cyanide  and  lead  oxide,  potassium 
cyanate  is  formed  : 

KCN  +  PbO  =  KCNO  +  Pb. 

Cyanic  acid  is  a  volatile,  acrid,  unstable  liquid.  It 
breaks  down  at  once  into  carbon  dioxide  and  ammonia 
in  presence  of  water  : 

COKE  +  H3O  =  NH3  +  C0a. 

The  potassium  salt  is  easily  soluble  in  water,  but  is 
decomposed  by  it,  yielding  ammonia  and  acid  potassium 
carbonate  : 

CONK  +  2H2O  =  KHC03  +  NH3. 

These  decompositions  of  cyanic  acid  and  the  cyanates 
further  exemplify  the  tendency  of  cyanogen  compounds 
to  undergo  decomposition  in  presence  of  water. 

Carbon  Bisulphide,  CS2.  —  Just  as  carbon  combines  di- 
rectly with  oxygen  to  form  the  dioxide,  so  it  combines 
directly  with  sulphur  to  form  the  disulphide  ;  but  there 
is  a  great  difference  in  the  ease  with  which  carbon  com- 
bines with  the  two  elements.  In  order  to  effect  combina- 
tion with  sulphur  a  very  high  temperature  is  necessary. 
The  compound  is  prepared  on  the  large  scale  by  heating 
charcoal  to  a  high  temperature  in  an  upright  cast-iron 
cylinder,  and  adding  sulphur  in  such  a  way  that  it  enters 
the  bottom  of  the  cylinder.  The  product  is  passed 


CARBON  BISULPHIDE.  405 

through  a  series   of   tubes   arranged   so   as   to   secure 
condensation. 

Carbon  disulphide  is  a  clear  liquid  which  has  a  high 
refractive  power.  It  boils  at  46°.  2.  When  pure  it  has 
a  pleasant  odor,  but  if  kept  for  a  time,  particularly  if 
water  is  present  in  the  vessel,  it  undergoes  slight  decom- 
position, and  products  of  extremely  disagreeable  odor 
are  formed.  It  can  generally  be  freed  from  these  by 
shaking  the  liquid  with  a  little  mercury  and  then  redis- 
tilling. It  burns  readily,  forming  carbon  dioxide  and 
sulphur  dixoide : 


CS,  +  302  =  C02  +  2S02 . 

In  nitric  oxide  it  burns  with  an  intensely  brilliant  flame, 
as  can  be  shown  by  filling  a  cylinder  with  the  gas,  adding 
a  few  drops  of  the  disulphide,  shaking  and  then  apply- 
ing a  flame.  A  column  of  brilliant  flame  rises  from  the 
mouth  of  the  cylinder  for  an  instant.  A  lamp  has  been 
constructed  in  which  this  flame  is  utilized.  It  is  of 
special  value  in  photographic  work. 

Carbon  disulphide  is  only  very  slightly  soluble  in 
water,  and  is  decomposed  by  it  only  very  slowly.  The 
disulphide  is  an  excellent  solvent  for  many  substances 
which  are  not  soluble  in  water,  as,  for  example,  fats, 
resins,  iodine,  and  one  of  the  modifications  of  sulphur 
and  of  phosphorus.  The  solution  of  iodine  in  it  has  a 
beautiful  violet  color ;  and  when  a  water  solution  con- 
taining a  little  free  iodine  is  shaken  with  carbon  disul- 
phide the  latter  acquires  a  violet  color  and  separates 
below  the  water. 

When  the  vapors  of  carbon  disulphide  and  hydrogen 
sulphide  are  passed  together  over  heated  copper,  methane 
and  cuprous  sulphide  are  formed,  as  has  been  stated. 
Methane  is  also  formed  when  the  vapors  of  carbon  disul- 
phide and  of  water  are  passed  over  ignited  iron.  While 
the  disulphide  is  not  easily  decomposed  by  water  at  the 
ordinary  temperature,  the  two  compounds  react  when 


406  INORGANIC  CHEMISTRY. 

heated  in  a  sealed  tube  to  150°,  the  products  being  car* 
bon  dioxide  and  hydrogen  sulphide  : 

CS2  +  2H2O  =  CO2  +  2H2S. 

Carbon  disulphide  finds  extensive  application  as  a 
solvent,  and  it  is  also  used  for  the  purpose  of  destroying 
phylloxera,  the  insect  which  is  so  destructive  to  grape- 
vines, particularly  in  the  wine  districts  of  France. 

Sulphocarbonic  Acid,  Thio-carbonic  Acid,  H2CS3. — Salts 
of  this  acid  are  formed  by  dissolving  carbon  disulphide 
in  concentrated  solutions  of  the  hydrosulphides.  Thus, 
when  it  is  dissolved  in  a  solution  of  sodium  hydrosulphide 
this  reaction  takes  place  : 

CS2  +  2NaSH  =  Na2CS3  +  H2S. 

The  reaction,  as  will  be  seen,  is  perfectly  analogous  to 
that  of  carbon  dioxide  upon  a  solution  of  sodium 
hydroxide : 

CO2  +  2NaOH  =  Na2CO3  +  H2O. 

The  salts  of  sulphocarbonic  acid  are  easily  decom- 
posed by  water  if  the  temperature  is  slightly  elevated, 
the  products  being  the  corresponding  carbonates  and 
hydrogen  sulphide : 

Na.CS,  +  3H2O  =  Na.CC),  +  3H2S. 

When  a  sulphocarbonate  is  treated  with  cold  dilute 
hydrochloric  acid,  the  free  acid  separates  as  a  dark  yel- 
low oil  of  a  very  disagreeable  odor.  This  readily 
undergoes  decomposition  into  carbon  disulphide  and 
hydrogen  sulphide : 

H2GS3  =•  CS2  -4-  H2S. 

This  reaction  is  again  perfectly  analogous  to  the  decom- 
position of  ordinary  carbonic  acid  into  carbon  dioxide 
and  water. 

Oxysulphides. — Products  intermediate  between  car- 
bonic acid  and  sulphocarbonic  acid  are  possible.  Such, 


CONSTITUTION  OF  CYANOGEN.  407 

for  example,  are  the  compounds  represented  by  the  for- 
mulas CO  •<  Q     ,  CS  -j          ,  etc. 


Sulphocyanic  Acid,  HCNS.  —  Just  as  the  cyanides  take 
up  oxygen  and  are  converted  into  cyanates,  so  also  they 
take  up  sulphur  and  are  converted  into  sulphocyanates  : 

KCN  +  S  =  KCNS. 

By  suspending  in  water  a  salt  of  the  acid,  the  metal  of 
which  forms  an  insoluble  sulphide  with  hydrogen  sul- 
phide, and  passing  this  gas  through  the  liquid,  a 
solution  of  the  acid  is  obtained.  When  the  solution  is 
boiled  the  acid  passes  over  partly  unchanged,  though  a 
part  is  decomposed  by  the  water  into  carbon  dioxide, 
carbon  disulphide,  and  ammonia  : 

2HCNS  +  2H2O  =  CO2  +  CS3  +  2NH3. 

The  ammonium  salt  of  sulphocyanic  acid  is  formed  by 
dissolving  carbon  disulphide  in  an  alcoholic  solution  of 
ammonia  : 

CS2  +  4NH3  =  (NH4)CNS  +  (NH4)2S. 

Constitution  of  Cyanogen  and  its  Simpler  Compounds.  — 
The  compounds  of  cyanogen  show,  in  general,  a  remark- 
able similarity  to  the  compounds  of  the  chlorine  group. 
The  hydrogen  compound  is  a  monobasic  acid  and  forms 
a  series  of  salts,  the  cyanides,  which  in  general  are  ana- 
logous to  the  chlorides.  Comparing  the  cyanides  with 
the  chlorides  it  is  clear  that  in  the  former  the  group  (CN), 
or  the  cyanogen  group,  plays  the  same  part  that  the  atom 
chlorine  plays  in  the  chlorides  : 

H(ON)  HC1 
K(CN)  KC1 
Hg(CN)a  HgCl, 

So  also  cyanic  acid  and  hypochlorous  acid  are  analo- 
gous : 

HO(CN)        HOC1. 


408  INORGANIC  CHEMISTRY. 

This  relation  suggests  that  which  is  observed  between 
the  ammonium  compounds  and  those  of  potassium  and 
sodium.  The  cyanogen  group  is  evidently  univalent,  as 
it  combines  with  one  atom  of  hydrogen,  one  of  potassium, 
etc.,  and  there  are  two  ways  in  which  we  can  conceive 
the  atoms  carbon  and  nitrogen  combined  to  form  a  uni- 
valent group.  If  the  nitrogen  is  trivalent  and  the  carbon 
quadrivalent  the  structure  is  that  represented  by  the 
formula  -C=N.  If,  on  the  other  hand,  the  nitrogen  is 
quinquivalent  and  the  carbon  quadrivalent  the  structure 
is  C=N-.  By  combination  of  the  first  group  with  hydro- 
gen a  compound  of  the  structure  H-C=N  would  result 
while  with  the  second  group  the  structure  of  the  hydro- 
gen compound  would  be  C=N-H,  In  the  one  case  the 
hydrogen  is  in  combination  with  carbon,  in  the  other  with 
nitrogen.  It  appears  probable  that  in  ordinary  hydro- 
cyanic acid  the  hydrogen  is  in  combination  with  carbon, 
the  structure  being  H-C=N.  This  is  in  accordance  with 
the  formation  of  the  acid  by  the  action  of  ammonia  upon 
chloroform,  which  is  most  readily  understood  on  the  as- 
sumption that  the  three  atoms  of  chlorine  are  replaced 
by  an  atom  of  nitrogen.  It  has  not  been  positively  de- 
termined which  of  the  two  possible  structures  above 
given  the  cyanogen  group  has  in  cyanic  acid.  In  one 
case  the  acid  would  have  the  structure  N=C— OH  ;  in  the 
other  it  would  be  C=N-OH.  It  may  also  be  O=C=NH. 
There  are  some  reasons  for  believing  that  the  ordinary 
cyanates  are  derived  from  an  acid  of  the  structure  rep- 
resented by  the  last  formula. 


CHAPTER  XXII. 

ELEMENTS  OF  FAMILY  IV,   GROUP  A: 
SILICON— TITANIUM— ZIRCONIUM— CERIUM— THORIUM. 

General. — While  the  elements  of  this  group  in  some 
respects  exhibit  resemblances  to  carbon  and  bear  to  it 
relations  similar  to  those  which  the  members  of  the 
chlorine  group  bear  to  fluorine,  the  members  of  the  sul- 
phur group  to  oxygen,  and  the  members  of  the  phos- 
phorus group  to  nitrogen,  yet  between  them  and  carbon 
there  are  some  remarkable  differences.  All  the  members 
of  the  group  except  titanium  combine  with  hydrogen. 
The  compounds  formed  have  the  formulas  SiH4,  ZrH2, 
CeH2,  and  ThH2.  The  power  to  form  homologous  series 
which  is  so  characteristic  of  carbon  is  entirely  wanting 
in  the  other  members  of  the  group.  With  the  members 
of  the  chlorine  group  they  all  form  compounds  analogous 
to  carbon  tetrachloride,  examples  of  which  are : 

SiCl4,        TiCl4,        ZrCl4,        CeF4,         ThCl4. 

Compounds  analogous  to  hexa-chlor-ethane,  C2Clg,  to 
tetra-chlor-ethylene,  C2C14,  and  to  octo-chlor-propane, 
C8Cle,are: 

Si2Cl8         Si2Cl4         Si3Cl8 
Ti2Cl6        Ti,Cl4 

All  the  elements  of  the  group  form  oxygen  compounds 
analogous  to  carbon  dioxide.  They  are  : 

Si02,        Ti02,        Zr02,        Ce02,        ThOa. 

The  first  three  are  acidic,  and  form  salts  which  in  com- 
position are  analogous  to  the  carbonates.  These  are  the 

(409) 


410  INORGANIC  CHEMISTRY. 

silicates,  titanates,  and  zirconates  of  the  general  formulas 

M3SiO3,        MaTiO3,        M3ZrO8. 

Cerium  and  thorium  oxides  are  basic.  These  facts 
suggest  the  relations  between  the  members  of  the  phos- 
phorus group.  The  oxides  of  the  last  two  members, 
antimony  and  bismuth,  are  basic,  although  the  oxide 
of  antimony  is  also  acidic  in  its  conduct  towards  the 
stronger  bases. 

The  compounds  of  silicon  are  very  abundant  in  nature  ; 
those  of  the  other  members  of  the  group  are  rare. 

SILICON,  Si  (At.  Wt.  28.18). 

Occurrence.— We  have  already  seen  what  an  exceed- 
ingly important  part  carbon  plays  in  animate  nature.  It 
is  interesting  to  note  that  silicon,  which  in  some  respects 
from  a  chemical  point  of  view  resembles  carbon,  is  one 
of  the  most  important  constituents  of  the  mineral  or  in- 
organic parts  of  the  earth.  It  occurs  chiefly  in  the  form 
of  the  dioxide,  SiO2,  commonly  called  silica,  or  silicon 
dioxide  ;  and  in  combination  with  oxygen  and  several  of 
the  common  metallic  elements,  particularly  with  sodium, 
potassium,  aluminium,  and  calcium,  in  the  form  of  the 
silicates.  Next  to  oxygen,  silicon  is  the  most  abundant  ele- 
ment in  the  earth.  There  are  extensive  mountain-ranges 
consisting  almost  entirely  of  the  dioxide,  SiO2,  in  the  form 
known  as  quartz  or  qiiartzite.  Other  ranges  are  made  up 
of  silicates,  which  are  compounds  formed  by  the  com- 
bination of  silicon  dioxide  and  bases.  The  clay  of 
the  vallevs  and  river-beds  also  contains  silicon  in 
large  quantity,  while  the  sand  found  so  abundantly 
on  the  deserts  and  at  the  sea-shore  is  largely  silicon 
dioxide. 

Preparation. — Silicon  does  not  occur  in  nature  in  the 
free  state.  The  oxide,  SiOa,  which  is  most  abundant  in 
the  form  of  sand,  is  decomposed  by  heating  it  with  potas- 
sium or  magnesium,  and  silicon  is  thus  set  free.  When 
magnesium  is  used  the  action  is  violent,  and  besides  the 
silicon  a  compound  of  silicon  and  magnesium  is  formed. 


SILICON.  411 

Silicon  has  also  been  made  by  heating  the  oxide  and 
carbon  in  the  electric  furnace,  and  by  decomposing  the 
chloride  with  potassium  : 

SiCl4  +  4K  =  Si  +  4KC1. 

The  best  way  to  make  it  is  by  heating  together  potassium 
fluosilicate,  K2SiFt,  sodium,  and  zinc: 

K2SiF6  +  4Na  =  4NaF  +  2KF  -f  Si. 

At  the  same  time  the  zinc  melts  and  the  silicon  which 
separates  dissolves  in  the  molten  zinc.  On  cooling,  it  is 
deposited  from  the  solution  in  beautiful  needle-shaped 
crystals,  around  which  the  zinc  solidifies  at  a  lower  tem- 
perature. By  treating  the  mass  with  hydrochloric  acid 
the  zinc  is  dissolved  and  the  crystals  of  silicon  are  left 
behind.  When  obtained  by  reduction  of  the  oxide  or  the 
chloride  by  means  of  potassium,  it  is  a  brown  amorphous 
powder.  If  made  by  decomposition  of  potassium  fluo- 
silicate by  aluminium,  it  is  deposited  from  the  molten 
aluminium  in  crystals  somewhat  resembling  graphite. 
Just  as  there  are  three  forms  of  carbon,  the  amorphous, 
graphite,  and  diamond,  so  there  are  three  corresponding 
forms  of  silicon,  the  amorphous  brown  powder,  the 
graphitoidal,  and  the  needles.  The  amorphous  variety 
is  converted  into  crystallized  silicon  by  continued  heat- 
ing at  a  high  temperature. 

Amorphous  silicon  acts  upon  hydrofluoric  acid,  form- 
ing silicon  tetrafluoride,  SiF4,  and  setting  hydrogen  free  : 


In  this  reaction  it  exhibits  one  of  the  properties  of  a 
base-forming  element.  Towards  other  acids,  however,  it 
is  indifferent.  It  is  not  acted  upon  by  sulphuric  acid, 
nor  by  nitric  acid,  nor  aqua  regia.  It  dissolves,  however, 
in  potassium  hydroxide,  forming  potassium  silicate,  in 
this  case  acting  like  an  acid-forming  element  : 

Si  +  2KOH  +  H,O  =  K2SiO3  +  2Hq. 


412  INORGANIC  CHEMISTRY. 

This  form  of  silicon  also  burns  in  the  air,  forming  the 
dioxide. 

Crystallized  silicon,  on  the  other  hand,  does  not  burn 
in  oxygen  at  the  highest  temperatures.  It,  however,  re- 
duces carbon  dioxide  and  decomposes  carbonates  at  a 
high  temperature.  It  is  also  oxidized  by  a  melting  mix- 
ture of  potassium  nitrate  and  the  hydroxide  or  carbonate. 
It  combines  with  nitrogen  at  a  high  temperature. 

Both  the  graphitoidal  and  needle-formed  crystals  of 
silicon  consist  of  regular  octahedrons.  Both  forms  have 
a  blackish-gray  color  and  a  metallic  lustre. 

Silicon  Hydride,  SiH4. — This  gas  is  obtained  mixed 
with  hydrogen  when  a  compound  of  magnesium  and  sili- 
con is  treated  with  hydrochloric  acid  : 

Mg2Si  +  4HC1  =  SiH4  +  2MgCl2. 

Thus  made,  it  takes  fire  when  it  comes  in  contact  with 
the  air,  and  the  act  is  accompanied  by  explosion.  The 
products  of  its  combustion  are  silicon  dioxide  and  water. 
When  pure  it  forms  a  colorless  gas  which  does  not  take 
fire  spontaneously  in  the  air  at  the  ordinary  temperature. 
If  it  is  diluted  with  hydrogen,  or  if  it  is  heated,  it  does 
take  fire.  When  burned  in  a  cylinder  or  narrow  tube, 
so  that  free  access  of  air  is  not  possible,  amorphous  sili- 
con is  deposited  upon  the  walls  of  the  vessel. 

Titanium,  Ti  (At.  Wt.  47.79). — Titanium  occurs  in  nature 
as  titanium  dioxide,  TiO2,  in  three  distinct  forms,  known 
as  rutile,  brookite,  and  anatase  ;  in  combination  with  iron, 
as  titaniferous  iron  which  contains  ferrous  titanate, 
JFeTiO3  ;  and  in  a  number  of  iron  ores  and  rare  minerals. 
The  element  is  obtained  in  the  free  state  by  decomposing 
potassium  fluotitanate,  K2TiF6,  with  potassium,  just  as 
silicon  is  obtained  by  decomposing  potassium  fluosilicate, 
.  K2SiF6,  with  potassium  or  sodium.  It  burns  when  heated 
in  the  air.  It  acts  upon  water  at  100°,  causing  the  evolu- 
tion of  hydrogen.  It  is  dissolved  by  hydrochloric  acid, 
forming  the  chloride,  Ti2Cl6.  At  a  high  temperature  it 
unites  directly  with  nitrogen  as  silicon  does.  Titanium 
does  not  form  a  compound  with  hydrogen. 


SILICON  TETRACHLOR1DE.  413 

Zirconium,  Zr  (At.  Wt.  89.72).— The  principal  form  in 
which  zirconium  occurs  in  nature  is  as  zircon,  which  is  a 
silicate  of  the  formula  ZrSiO4,  derived  from  normal 
silicic  acid,  Si(OH)4,  by  the  replacement  of  the  four  hy- 
drogen atoms  by  a  quadrivalent  atom  of  zirconium.  The 
element  is  obtained  in  the  free  condition  by  decomposing 
potassium  fluozirconate  by  heating  it  with  aluminium  to 
a  high  temperature.  In  this  way  it  is  obtained  in  crystal- 
lized form,  somewhat  resembling  antimony.  It  does  not 
burn  in  the  air.  It  is  dissolved  by  hot  concentrated  hy- 
drochloric acid,  and  when  heated  in  a  current  of  hydro- 
chloric acid  gas.  The  product  is  the  tetrachloride,  ZrCl4 ; 
and  the  same  compound  is  formed  when  chlorine  acts 
directly  upon  zirconium. 

Thorium,  Th.  (At.  Wt.  230.87). — This  element  occurs 
principally  in  the  mineral  thorite,  which  is  essentially  a 
silicate  of  thorium,  ThSiO4 ,  analogous  to  zircon.  It  is 
obtained  free  by  treating  the  chloride  with  silicon  or 
potassium.  At  high  temperatures  it  burns  in  the  air, 
forming  thorium  dioxide,  ThO2. 

Cerium  so  much  resembles  the  two  elements  lanthanum 
and  didymium,  that  although  it  falls  in  the  same  group  as 
silicon,  and  resembles  the  elements  of  this  group  in  some 
respects,  it  seems  advisable  to  postpone  its  study  until 
lanthanum  and  didymium  are  taken  up. 

COMPOUNDS  OF  THE  ELEMENTS  OF  THE  SILICON  GKOUP 

WITH  THOSE  OF  THE  CHLOKINE  GROUP. 

Silicon  Tetrachloride,  SiCl4.— This  compound  is  formed 
when  silicon  is  heated  in  a  current  of  chlorine,  and  by 
passing  a  current  of  dry  chlorine  over  a  heated  mixture 
of  silicon  dioxide  and  carbon.  Under  these  latter  cir- 
cumstances the  following  reaction  takes  place  : 

SiO,  +  20  +  2C1,  =  SiCl4  +  2CO. 

Carbon  acting  alone  upon  silicon  dioxide  cannot  reduce 
it,  nor  has  chlorine  acting  alone  the  power  to  convert  it 
into  the  chloride.  When,  however,  carbon  and  chlorine 


414  INORGANIC  CHEMISTRY. 

act  together  both  reactions  take  place.  The  tetrachlo- 
ride  is  a  colorless  liquid.  It  is  decomposed  by  water, 
forming  silicic  acid  and  hydrochloric  acid.  The  reaction 
probably  takes  place  as  represented  in  the  following 
equation : 

SiCl4  +  4H20  =  Si(OH)4  +  4HC1. 

The  normal  acid  thus  formed  breaks  down  very  readily, 
however,  forming  the  ordinary  acid  of  the  formula 
SiO(OH)2  or  H2SiO3,  corresponding  to  carbonic  acid, 
H2CO, 

Silicon  Hexachloride,  Si2Cl6,  is  formed  when  silicon 
tetrachloride  is  heated  with  silicon  : 

3SiCl4  +  Si  =  2Si2Cle. 

When  heated  to  a  sufficiently  high  temperature  it  is  de- 
composed, yielding  silicon  and  the  tetrachloride  : 

2Si2Cl6  =  3SiCl4  +  Si. 

"Water  decomposes  it,  forming  the  corresponding  hy- 
droxyl  derivative,  which  loses  water  and  forms  the  acid 
Si202(OH)2 : 

Si2Cl6  +  6H2O  =  Si2(OH)6  +  6HC1. 
Si2(OH)6  =  Si202(OH)2  +  2H20. 

The  product  is  a  disilicic  acid,  in  some  respects  analo- 
gous to  disulphuric  acid. 

Similar  compounds  of  silicon  with  bromine  and  iodine 
are  known. 

Silicon  Tetrafluoride,  SiF4. — This  is  one  of  the  most 
interesting  of  the  compounds  which  silicon  forms  with 
the  members  of  Family  VII.  It  is  made  by  treating 
silicon  dioxide  with  hydrofluoric  acid.  This  action  is 
secured  by  treating  a  mixture  of  silicon  dioxide  (sand) 
and  calcium  fluoride  (fluor-spar)  with  concentrated  sul- 
phuric acid,  when  two  reactions  take  place  : 

CaF2  +  H2S04  =  CaSO4  +  2HF  ; 
Si02  +  4HF     =  2H20   +  SiF4 . 


FLUOSILICIC  ACID.  415 

The  tetrafluoride  escapes  as  a  colorless  gas,  which  forms 
thick  clouds  in  moist  air  on  account  of  the  action  of 
water  upon  it. 

Water  decomposes  the  tetrafluoride,  as  it  does  the 
tetrachloride.  The  first  action  probably  consists  in  the 
formation  of  normal  silicic  acid  and  hydrofluoric  acid,  the 
normal  acid  then  breaking  down  by  loss  of  water  and 
yielding  the  ordinary  form  of  silicic  acid : 

SiF4  +  4H2O  =  Si(OH)4  +  4HF  ; 
Si(OH)4  =  SiO(OH)2  +  H2O. 

The  silicic  acid  thus  formed  separates  as  a  gelatinous 
mass.  At  the  same  time  the  hydrofluoric  acid  acts  upon 
some  of  the  silicon  tetrafluoride,  forming  the  compound 
fluosilicic  acid,  which  has  the  formula  H2SiF6 : 

SiF4  +  2HF  =  H2SiF6. 

The  complete  action  may  be  represented  in  one  equa- 
tion, as  follows : 

3SiF4  +  3H2O  =  H2Si03  +  2H2SiF6. 

The  fluosilicic  acid  remains  in  solution  in  the  water,  and 
by  treating  this  solution  with  carbonates  or  hydroxides 
of  the  metallic  elements  the  salts  known  as  the  fluosili- 
cates  are  obtained.  The  solution  of  the  acid  can  be  con- 
centrated to  a  certain  extent  in  a  platinum  vessel,  but  it 
breaks  down  into  silicon  tetrafluoride  and  hydrofluoric 
acid  when  it  becomes  concentrated.  If  more  potassium 
hydroxide  than  is  required  to  neutralize  the  acid  is  added 
to  the  solution,  decomposition  ensues,  with  formation  of 
silicic  acid : 

H2SiF6  +  6KOH  =  6KF  +  H2SiO3  +  3H2O. 

By  water  alone,  however,  the  acid  is  not  decomposed, 
and  the  salts  are  fairly  stable.  When  heated,  the  salts 
give  off  silicon  tetrafluoride,  and  fluorides  are  left  behind  •• 

K2SiF6  =  2KF  +  SiF4. 


416  INORGANIC  CHEMISTRY. 

Constitution  of  Pluosilicic  Acid.  —  Attention  has  already 
been  called  to  the  fact  that  fluosilicic  acid  and  silicic  acid 
seem  to  be  analogous  substances,  and  that  the  former 
may  be  regarded  as  derived  from  the  latter  by  the  sub- 
stitution of  six  fluorine  atoms  for  the  three  oxygen 
atoms.  According  to  this,  fluorine  has  a  valence  higher 
than  one,  and  this  accords  with  the  fact  that  at  the  or- 
dinary temperature  the  density  of  hydrofluoric  acid  is 
greater  than  that  required  by  the  formula,  HF.  Assum- 
ing, then,  that  fluorine  may  act  as  a  bivalent  or  a  tri- 
valent  element,  and  for  the  present  purpose  it  is  imma- 
terial which  view  is  taken,  the  relation  between  silicic 
acid  and  fluosilicic  acid  is  shown  by  the  following  for- 
mulas : 


Silicic  acid  Fluosilicic  acid 

It  is  commonly  held  that  the  acid  is  a  "  double  com- 
pound "  made  by  the  union  of  one  molecule  of  silicon  tetra- 
fluoride  with  two  molecules  of  hydrofluoric  acid,  and  rep- 
resented by  the  formula  SiF4.2HF.  This  is  not  even  an 
attempt  at  an  explanation  of  the  fact  that  the  composition 
of  the  acid  is  so  similar  to  that  of  silicic  acid.  The  above 
explanation  is,  however,  in  accordance  with  the  composi- 
tion of  a  large  number  of  similar  "  double  compounds," 
in  which  not  only  fluorine,  but  chlorine,  bromine,  and 
iodine  enter. 

Titanium  Tetrachloride,  TiCl4,  is  formed  by  the  direct 
action  of  chlorine  on  titanium,  and  also  by  passing  dry 
chlorine  over  a  mixture  of  carbon  and  titanium  dioxide. 
It  is,  like  silicon  tetrachloride,  a  liquid.  It  forms  crys- 
tallized compounds  with  water.  When  heated  with 
water  it  is  decomposed,  yielding  titanium  dioxide  and 
hydrochloric  acid  : 

TiCl4  +  2H,0  =  TiO,  +  4HC1. 

Titanium  also  forms  with  chlorine  the  compounds  Ti2Cl9 
and  Ti2Cl4. 


THORIUM  TETRAFLUORIDE.  41? 

Titanium  Tetraihioride,  TiF4,  is  formed  in  the  same 
way  as  silicon  tetrafluoride,  by  treating  a  mixture  of 
titanium  dioxide  and  fluor-spar  with  concentrated  sul- 
phuric acid,  and  by  dissolving  titanium  dioxide  in  hydro- 
fluoric acid.  When  treated  with  water  it  forms  a  com- 
pound analogous  to  fluosilicic  acid,  called  fluotitanic 
acid,  H2TiF6,  which  yields  well  characterized  salts,  the 
fluotitanates. 

Zirconium  Tetrachloride,  ZrCl4,  is  not  completely  decom- 
posed by  water,  only  half  the  chlorine  being  replaced  by 
oxygen,  forming  a  product,  zirconium  oxychloride,  ZrOCl2 : 

ZrCl4  +  H20  =  ZrOCl2  +  2HC1. 

This  is  in  accordance  with  the  fact  that  zirconium  acts 
as  a  base-forming  as  well  as  an  acid-forming  element. 
The  chlorides  of  silicon  and  titanium  are  completely  de- 
composed by  water,  as  they  are  acid-forming. 

The  tetrafluoride  of  zirconium  is  obtained  from  zircon 
or  zirconium  silicate  by  mixing  the  finely  powdered 
mineral  with  fluor-spar  and  passing  hydrochloric  acid 
gas  over  it  at  a  high  temperature  : 

ZrSi04  +  2CaF2  +  2HC1  =  CaCl2  +  CaSiO3  +ZrF4  +  H2O. 

With  metallic  fluorides  the  tetrafluoride  forms  salts  of 
fluozirconic  acid,  H2ZrF6,  analogous  to  fluosilicic  and  fluo- 
titanic acids. 

Thorium  Tetrachloride,  ThCl4,  is  not  decomposed  by 
water  at  the  ordinary  temperature,  but  if  its  solution 
is  evaporated  to  dryness  hydrochloric  acid  is  given  off 
and  thorium  dioxide  is  left : 

ThCl4  +  2H20  =  Th02  +  4HC1. 

Thorium  Tetrafluoride,  ThF4,  is  easily  made  by  treat- 
ing the  tetrachloride  with  hydrofluoric  acid.  With  po- 
tassium fluoride  it  forms  a  salt  of  the  formula  K2ThF6, 
or  potassium  fluothorate.  The  chloride  also  forms  a  simi- 
lar salt  with  potassium  chloride,  potassium  chlorthorate, 
K2ThCl6. 


418  INORGANIC  CHEMISTRY. 

Comparison  of  the  Chlorides  of  Family  IV  with  those 
of  Family  V. — In  studying  the  chlorides  formed  by  the 
members  of  the  phosphorus  group  it  was  found  that  the 
chlorides  of  phosphorus  are  readily  decomposed  by  water, 
forming  the  corresponding  acids,  and  that  the  same  is 
true  of  the  chloride  of  arsenic ;  but  that  the  trichlorides 
of  antimony  and  bismuth  are  only  partly  decomposed 
by  water,  yielding  oxychlorides.  In  the  silicon  group 
we  find  now  similar  differences  between  the  members 
with  low  atomic  weights  and  those  with  high  atomic 
weights.  The  chlorides  of  silicon  and  titanium  are  com- 
pletely decomposed  by  water  at  the  ordinary  tempera- 
ture, while  that  of  zirconium  is  only  half  decomposed, 
and  that  of  thorium  is  not  decomposed  except  at  high 
temperature. 

COMPOUNDS  OF  THE  MEMBERS  or  THE  SILICON  GROUP 
WITH  OXYGEN,  AND  WITH  OXYGEN  AND  HYDROGEN. 

Silicon  Dioxide,  SiO2. — This  compound  occurs  very 
abundantly  in  nature  in  many  different  forms,  both  crys- 
tallized and  amorphous.  Quartz  is  a  form  of  crystallized 
silicon  dioxide  which  is  found  very  widely  distributed. 
It  crystallizes  in  the  hexagonal  system  in  prisms  and 
pyramids,  the  crystals  sometimes  attaining  great  size 
and  beauty.  Another  form  of  the  crystallized  compound 
is  that  known  as  tridymite.  Like  quartz  it  crystallizes  in 
the  hexagonal  system,  but  the  characteristic  forms  are 
not  the  same  as  those  of  quartz.  Further,  it  nearly 
always  occurs  in  triplet  crystals.  The  finer  crystals  of 
quartz  are  generally  called  rock-crystal ;  the  crystalline 
variety  in  which  the  crystals  are  not  well  developed  is 
called  quartzite.  The  amorphous  varieties  of  silicon  di- 
oxide frequently  contain  water  in  combination,  or,  rather, 
they  are  hydroxides  of  silicon.  Examples  of  these  forms 
are  opal,  agate,  amethyst,  carnelian,  flint,  sand,  chalced- 
ony. Some  of  these  are  colored  by  small  quantities  of 
other  substances  contained  in  them.  Carnelian  owes 
its  color  to  a  compound  of  iron,  probably  ferric  oxide ; 
flint  contains  small  quantities  of  organic  matter.  The 


SILICON  DIOXIDE.  419 

specific  gravity  of  the  crystallized  varieties  is  higher 
than  that  of  the  amorphous  varieties,  and  there  are  also 
some  chemical  differences  between  them  which  will  be 
referred  to  below. 

Pure  silicon  dioxide  can  be  made  by  melting  sand  or 
a  finely  powdered  silicate  with  sodium  carbonate,  when 
sodium  silicate  is  formed.  This  is  soluble  in  water,  and 
when  hydrochloric  acid  is  added  to  the  solution  silicic 
acid  separates  in  the  form  of  a  gelatinous  mass.  By 
evaporating  the  mass  to  complete  dryness,  moistening 
with  concentrated  hydrochloric  acid,  and  after  a  time 
treating  with  water,  everything  dissolves  except  silicon 
dioxide,  which  is  perfectly  pure  and  in  a  very  finely  di- 
vided state.  It  can  also  be  obtained  pure  by  passing  sili- 
con tetrafluoride  into  water.  As  we  have  seen,  a  form 
of  silicic  acid  separates  under  these  circumstances. 
This,  when  filtered,  dried,  and  ignited,  yields  perfectly 
pure  silicon  dioxide. 

Properties.  —  Silicon  dioxide  is  insoluble  in  water  and 
in  most  acids.  It  dissolves,  however,  in  hydrofluoric 
acid,  forming  the  tetrafluoride.  It  requires  the  tempera- 
ture produced  by  the  oxyhydrogen  blow-pipe  to  melt  it. 
The  amorphous  varieties  are  more  easily  acted  upon  by 
other  substances  than  the  crystallized.  Thus,  hydro- 
fluoric acid  acts  much  more  readily  upon  them.  When 
the  amorphous  compound  is  boiled  with  solutions  of 
potassium  or  sodium  hydroxide,  or  of  the  carbonates  of 
these  metals,  it  dissolves,  forming  the  corresponding  sili- 
cate : 


2KOH  +  SiO2  =  K2SiO3  +  H2O. 

The  crystallized  varieties  are  not  dissolved  in  this  way. 
All  forms  of  the  dioxide  act  upon  melting  hydroxides  or 
carbonates  of  potassium  or  sodium,  and  form  the  corre- 
sponding silicates. 

Uses.  —  Plants  take  up  silicon  dioxide  from  the  soil, 
and  this  being  deposited  in  various  part  of  their  tissues, 
gives  them  the  necessary  firmness.  -  Straw,  for  example, 


420  INORGANIC  CHEMISTRY. 

is  rich  in  silicon  dioxide.  Horse-tail,  a  plant  of  the 
genus  Equisetum,  is  so  rich  in  finely  divided  silicon  di- 
oxide that  it  is  used  for  polishing.  There  are  great 
natural  deposits  of  finely  divided  silicon  dioxide  known 
as  infusorial  earth.  This  consists  of  the  remains  of  dia- 
toms. And  finally  silicon  dioxide  is  found  in  the  hair, 
in  feathers,  and  in  egg  albumen.  Silicon  dioxide  finds 
extensive  application  in  the  manufacture  of  mortar,  glass, 
and  porcelain.  Ordinary  glass,  as  we  shall  see,  is  a 
silicate  of  calcium  and  potassium  or  sodium,  which  is 
made  by  melting  together  sand  and  the  carbonates  of 
the  metals  mentioned. 

Silicic  Acid. — There  are  many  varieties  of  silicic  acid, 
all  of  which  can,  however,  be  referred  to  the  normal  acid, 
Si(OH)4.  This  normal  acid  is  not  known  in  the  free 
state  in  pure  condition,  but  it  is  probably  contained  in 
the  gelatinous  precipitate  which  is  formed  when  silicon 
tetrachloride  or  tetrafluoride  is  decomposed  by  water  : 

SiCl4  +  4H2O  =  Si(OH)4  +  4HC1. 

This  cannot,  however,  be  isolated,  as,  even  by  standing, 
it  loses  a  molecule  of  water,  and  passes  into  the  form 
H2Si03: 

Si(OH)4  =  OSi(OH)2  +  H2O. 

This  is  the  form  from  which  most  of  the  ordinary  sili- 
cates are  derived.  It  cannot  be  isolated  in  pure  con- 
dition, for  when  filtered  off  and  exposed  to  the  air  it 
loses  more  water,  and  when  heated  to  a  sufficiently  high 
temperature  it  is  converted  into  silicon  dioxide. 

OSi(OH),  =  Si0.2  +  H2O. 

When  potassium  or  sodium  silicate  in  solution  is  treated 
with  hydrochloric  acid,  most  of  the  silicic  acid  separates 
in  the  form  of  a  gelatinous  mass  if  the  solution  is  con- 
centrated. If,  however,  the  solution  is  dilute,  a  consid- 
erable part  of  the  acid  remains  in  solution.  Further,  if 
a  concentrated  solution  of  the  silicate  of  potassium  or 


SILICIC  ACID.  421 

sodium  is  poured  quickly  into  hydrochloric  acid,  or  if 
the  acid  is  poured  quickly  into  the  solution  of  the  sili- 
cate, the  silicic  acid  remains  in  solution.  If,  however,  the 
solutions  are  brought  together  drop  by  drop  the  silicic 
acid  separates.  From  these  solutions  of  silicic  acid  am- 
monia or  ammonium  carbonate  throws  down  the  acid. 

A  solution  of  pure  silicic  acid  can  be  obtained  by 
means  of  dialysis.  It  has  been  found  that  solutions  of 
different  substances  pass  with  different  degrees  of  ease 
through  porous  membranes,  just  as  gases  differ  as  re- 
gards the  ease  with  which  they  pass  through  porous  dia- 
phragms. This  fact  concerning  gases  was  referred  to  in 
connection  with  hydrogen.  Now,  while  some  solutions 
pass  readily  through  parchment  paper,  others  pass 
through  with  difficulty,  and  some  do  not  pass  through  at 
all.  A  dialyser,  or  an  apparatus  used  in  dialysis,  may  be 
made  by  tying  a  piece  of  parchment  paper  over  the  mouth 
of  a  ring-formed  glass  or  rubber  vessel,  and  placing  this  in 
another  shallow  vessel.  Pure  water  is  put  in  the  outer 
vessel,  and  the  solution  for  dialysis  in  the  inner  one. 
The  arrangement  is  illustrated  in  Fig.  13. 


FIG.  13. 

In  the  figure  aa  is  the  hoop  of  gutta-percha,  and  5  is  the 
parchment  paper.  When  now  the  solution  containing  hy- 
drochloric acid,  sodium  chloride,  and  silicic  acid  is  put  in 
the  dialyser,  the  hydrochloric  acid  and  sodium  chloride 
pass  readily  through  the  membrane,  while  the  silicic  acid 
is  left  behind,  and  in  the  course  of  a  few  days,  if  the  water 
in  the  outer  vessel  is  renewed,  the  solution  of  silicic 
acid  in  the  inner  vessel  will  be  found  to  be  free  from  the 
other  substances.  This  solution  can  be  evaporated  to 


422  INORGANIC  CHEMISTRY, 

some  extent  by  boiling,  but  when  a  certain  concentration 
is  reached  the  acid  separates.  In  a  vacuum  such  a  solu- 
tion can  be  evaporated  further  without  the  formation  of 
a  deposit.  Finally,  there  is  left  a  transparent  mass  which 
has  approximately  the  composition  represented  by  the 
formula  H2SiO3.  The  dialysed  solution  of  silicic  acid  is 
coagulated  by  a  very  dilute  solution  of  sodium  or  potas- 
sium carbonate,  and  by  carbon  dioxide  itself. 

When  the  solutions  containing  silicic  acid  are  evapo- 
rated to  complete  dryness  the  acid  is  converted  into  sili- 
con dioxide  and  other  insoluble  hydrates.  This  residue 
is  called  insoluble  silicic  acid.  When  this  is  treated  with 
hydrochloric  acid  and  water  it  remains  undissolved,  and 
if  filtered  off  and  ignited  it  leaves  a  residue  of  silicon  di- 
oxide. To  sum  up,  then  :  Whenever  silicic  acid  is  formed 
in  a  solution  it  is  a  more  or  less  complex  derivative  of 
normal  silicic  acid,  and  is  somewhat  soluble  in  water,  but 
by  the  processes  just  described  the  soluble  acid  is  con- 
verted into  insoluble  silicic  acid,  as  explained. 

Poly  silicic  Acids. — Silicic  acid  is  remarkable  for  the 
great  number  of  derivatives  which  it  yields.  Most  of 
these  bear  to  the  normal  acid  relations  similar  to  those 
which  the  various  forms  of  phosphoric  acid  bear  to  nor- 
mal phosphoric  acid,  and  the  various  forms  of  periodic 
acid  to  normal  periodic  acid.  It  has  already  been  stated 
that  salts  of  the  acid  H2SiO3  are  more  common  than 
those  of  the  normal  acid.  Among  the  salts  of  the  normal 
acid  are  zircon,  ZrSiO4,  and  thorite,  ThSiO4.  The  or- 
dinary silicates  of  potassium  and  sodium  are  derived 
from  the  acid  H2SiO3 ;  so  also  are  wollastonite,  CaSiO3, 
and  enstatite,  MgSiO3. 

Disilicic  Acid  is  derived  from  "ordinary  silicic  acid  by 
loss  of  one  molecule  of  water  from  two  molecules  of  the 
acid : 


Its  composition  is,  therefore,  H^Si^O,,  which   may  be 
W 


TRISIL1C1C  ACIDS.  423 

written  O3Si,(OH)2.  Another  form  of  disilicic  acid  is  de- 
rived from  two  molecules  of  the  normal  acid  by  loss  of 
one  molecule  of  water  : 

2Si(OH)4  =  OSi,(OH)6  +  H20. 

The  well-known  mineral  serpentine  is  apparently  the 
magnesium  salt  of  this  acid.  It  is  represented  by  the 
formula  Mg,Si9O,. 

Trisilicic  Acids  are  derived  from  three  molecules  of 
the  normal  acid  or  the  ordinary  acid  by  loss  of  different 
numbers  of  molecules  of  water.  Thus,  by  loss  of  two 
molecules  the  normal  acid  would  yield  a  product  HflSi3O10. 
By  loss  of  two  molecules  of  water  this  trisilicic  acid 
would  yield  an  acid  of  the  formula  H4Si3O8.  The  struc- 
ture of  the  first  acid  is  expressed  by  formula  I,  and  of 
the  second  by  formula  II,  below  given : 


Si 


Sij(OH), 


i   (OH),  Si 


OH  (OK 

O  Si  J0 


O  O 

O  Si]o 
O  O 

(OH),  Si|o8Al 
n.  in. 


Orthoclase  or  ordinary  feldspar  is  the  aluminium- 
potassium  salt  of  the  second  form  of  trisilicic  acid,  in 
which  one  atom  of  hydrogen  is  replaced  by  potassium, 
and  three  by  an  aluminium  atom,  as  shown  in  formula 
III  above, 

Titanium  Dioxide,  TiOa — As  has  been  stated,  this  is 
one  of  the  principal  forms  in  which  titanium  is  found  in 
nature.  There  are  three  natural  crystallized  varieties — 
rutile,  brookite,  and  anatase.  In  order  to  prepare  the 
pure  dioxide  from  one  of  the  natural  forms,  it  is  melted 
in  finely  powdered  condition  with  potassium  carbonate, 
when  it  is  converted  into  potassium  titanate,  K2TiO3,  the 
reaction  being  entirely  analogous  to  that  which  takes 
place  when  silicon  dioxide  is  treated  in  the  same  way : 

K2C03  +  Si02  =  K2Si03  +  C02 ; 
K2C03  +  TiO3  =  K2Ti03  +  C0a. 


424  INORGANIC  CHEMISTRY. 

When  titanic  acid  is  precipitated  from  a  solution  of  a 
titanate  it  appears  as  a  hydroxide,  the  composition  of 
which  varies  from  Ti(OH)4,  or  normal  titanic  acid,  to 
H2Ti2O5,  a  dititanic  acid.  When  these  substances  are 
ignited  they  yield  titanium  dioxide.  The  hydroxides  of 
titanium  conduct  themselves  somewhat  like  those  of  sili- 
con. They  are  to  some  extent  soluble  in  water,  and  when 
these  solutions  containing  sulphuric  acid  are  much 
diluted  and  boiled,  the  titanium  is  all  precipitated  as  a 
hydroxide.  Titanium  dioxide  forms  some  salts  with 
acids,  among  which  the  following  are  examples  Ti(SO4)a 
and  TiO(SO4).  The  former  is  normal  titanium  sulphate, 
the  latter  titanyl  sulphate,  in  which  the  bivalent  group, 
TiO,  or  titanyl,  takes  the  place  of  two  hydrogen  atoms. 

Zirconium  Dioxide,  ZrO2,  is  obtained  by  a  rather  com- 
plicated series  of  reactions  from  zircon.  It  dissolves  in 
molten  potassium  or  sodium  carbonate,  forming  the  cor- 
responding zirconate,  K2ZrO3.  The  sodium  salt  of  nor- 
mal zir conic  acid,  Na4ZrO4,  has  also  been  obtained. 

The  dioxide  forms  salts  with  acidrf,  among  which  two 
of  the  sulphates  are  of  special  interest.  One  has  the 
composition  ZrSO6,  and  the  other  Zr3(SO6)2.  The  former 
is  to  be  regarded  as  the  salt  of  the  acid  OS(OH)4,  formed 
by  substituting  one  atom  of  zirconium  for  the  four  atoms 
of  hydrogen ;  the  other  is  the  salt  of  normal  sulphuric 
acid,  S(OH)e,  formed  by  substituting  zirconium  for  all 
the  hydrogen. 

Thorium  Dioxide,  ThO2 ,  does  not  form  thorates  as  the 
dioxides  of  the  other  members  of  the  group.  It  does, 
however,  form  salts  with  acids.  In  these,  thorium  acts 
as  a  quadrivalent  element. 

Silicides  are  compounds  of  silicon  with  other  elements, 
as,  for  example,  with  carbon.  These  two  elements  com- 
bine forming  an  extremely  interesting  compound  carbon 
silicide,  CSi,  which  is  manufactured  on  the  large  scale  and 
known  in  the  market  as  carborundum.  This  is  made  by 
heating  a  mixture  of  quartz  sand,  coke,  and  common 
salt,  or  sodium  chloride,  in  the  electric  furnace  to  3500°,  • 
when  the  reaction  represented  below  takes  place : 

SiO,  +  20  =  CSi  +  2CO. 


FAMILY  IV,  GROUP  B.  425 

The  product  is  in  the  main  crystallized,  the  crystals 
being  bluish  or  yellowish-green.  They  have  the  specific 
gravity  3.22  to  3.12.  The  silicide  is  said  to  be  colorless 
when  perfectly  pure.  It  scratches  ruby  and  chrome- 
steel,  and  on  account  of  its  hardness  it  is  much  prized 
as  a  polishing  agent,  being  used  to  a  considerable  ex- 
tent in  place  of  emery.  Pure  carbon  silicide  is  insoluble 
in  nearly  all  ordinary  solvents,  including  hydrochloric, 
nitric,  sulphuric,  and  hydrofluoric  acids.  It  is,  however, 
decomposed  by  fusing  caustic  alkalies  or  their  car- 
bonates. 

Many  other  silicides  have  been  made,  several  of  which 
are  well  crystallized  compounds. 

FAMILY  IV,  GROUP  B. 

Allied  to  the  members  of  the  silicon  group,  yet  differ- 
ing from  them  in  some  important  particulars,  are  the 
three  elements  germanium,  tin,  and  lead.  Of  these  the 
first  two  are  more  acidic  in  character  than  the  last.  The*y 
combine  with  chlorine  in  two  proportions,  forming  the 
chlorides  GeCla,  SnCla,  PbCl,',  GeCl4,  SnCl4,  PbCl4.  With 
oxygen  they  unite,  forming  the  compounds  GeO2,  SnO2, 
and  PbO2.  Stannic  oxide,  SnO2,  and  lead  peroxide,  PbO2, 
form  salts  with  bases,  and  these  have  the  composition 
represented  by  the  general  formulas  M2SnO3  and  M2PbO3, 
and  are  therefore  analogous  to  the  silicates  and  titanates. 
On  the  other  hand,  further,  salts  are  known  which  are 
derived  from  the  oxide  PbO.  These  have  the  general 
formula  M2PbO2,  and  are  to  be  regarded  as  salts  of  an 
acid,  Pb(OH)2.  These  salts  are  not  stable,  and  are  not 
easily  obtained.  Most  of  the  derivatives  of  lead  are 
those  in  which  it  plays  the  part  of  a  base-forming  ele- 
ment. It  will  therefore  be  better  to  postpone  its  study 
until  it  is  taken  up  under  the  general  head  of  the  base- 
forming  elements.  Notwithstanding,  further,  the  marked 
analogy  between  some  of  the  compounds  of  tin  and  those 
of  the  members  of  the  silicon  group,  it  appears  on  the 
whole  advisable  to  treat  of  this  element  in  company  with 
lead,  which  it  also  resembles  in  many  respects. 


CHAPTER  XXIII. 

CHEMICAL  ACTION. 

Retrospective. — "We  have  been  studying  the  principal 
elements  of  four  families  and  the  compounds  which  they 
form  with  one  another.  No  matter  how  simple  or  how 
complex  the  chemical  changes  studied  were,  certain  fun- 
damental laws  governing  all  cases  of  chemical  action  were 
found  to  hold  good.  These  laws  have  been  discussed, 
but  it  will  be  well  to  recall  them  here  before  taking 
up  other  laws  which  are  intimately  connected  with  them. 
The  first  great  law  of  chemical  change  is 

I.  The  laiv  of  conservation  of  mass. 

*  According  to  this,  the  amount  of  matter  is  not  changed 
by  a  chemical  act. 
The  second  law  is 

II.  The  law  of  definite  proportions. 

According  to  this,  the  composition  of  every  compound 
is  always  the  same. 
The  third  law  is 

III.  The  law  of  multiple  proportions. 

According  to  this,  the  different  masses  of  any  element 
which  combine  with  a  fixed  mass  of  another  or  others 
bear  simple  relations  to  one  another. 

To  account  for  the  laws  of  definite  and  multiple  pro- 
portions the  Atomic  Theory  was  proposed. 

According  to  this,  each  element  is  made  up  of  particles 
of  definite  weight,  which  are  chemically  indivisible,  and 
chemical  action  consists  in  union  or  separation  of  these 
particles.  These  hypothetical  particles  are  called  atoms. 
The  elements  must  combine  in  the  proportion  of  their 
atomic  weights  or  of  simple  multiples  of  these,  if  the 
atomic  theory  is  true. 

Further  study  showed  that  it  is  necessary  to  assume 

(426) 


CHEMICAL  ACTION.  427 

the  existence  of  larger  particles  than  the  atoms,  viz.,  the 
molecules.  According  to  the  theory  of  molecules,  every 
chemical  compound  and  element  is  made  up  of  mole- 
cules, which  are  the  smallest  particles  having  the  same 
general  properties  as  the  mass.  These  molecules  are 
made  up  of  atoms  which,  in  the  case  of  compounds,  are 
of  different  kinds,  and  in  the  case  of  elements,  of  the 
same  kind.  In  the  case  of  a  few  elements  the  atom  ap- 
pears to  be  identical  with  the  molecule. 

From  the  study  of  gases  the  conclusion  is  reached  that 
in  equal  volumes  of  all  gases  under  standard  conditions 
there  is  always  the  same  number  of  molecules  (Avoga- 
dro's  law).  This  gives  us  a  means  of  determining  the 
relative  weights  of  molecules  of  gaseous  substances  ;  and 
from  these  molecular  weights  it  is  possible  to  draw  con- 
clusions in  regard  to  the  atomic  weights  of  those  elements 
which  enter  into  the  composition  of  the  compounds  thus 
studied. 

The  formulas  of  chemical  compounds  are  intended  to 
be  molecular  formulas.  They  are  intended  to  tell  of  what 
atoms  and  of  how  many  atoms  the  molecules  represented 
are  made  up. 

The  method  of  determining  molecular  weights  based 
upon  Avogadro's  law  is  applicable  only  to  gaseous  sub- 
stances, or  to  such  as  can  be  converted  into  gas  without 
undergoing  decomposition.  While  many  of  the  com- 
pounds with  which  we  have  had  to  deal  are  of  this  char- 
acter, many  of  them  are  not,  and  in  regard  to  the  mole- 
cular weights  of  these,  we  must  be  in  doubt  unless  some 
other  method  applicable  to  liquids  and  solids  is  avail- 
able. So,  too,  the  atomic  weights  of  those  elements  which 
enter  into  the  composition  of  gaseous  compounds  can  be 
deduced  from  the  molecular  weights,  but  plainly  those 
which  do  not  enter  into  the  composition  of  such  com- 
pounds demand  some  other  method.  For  determining 
the  atomic  weights  of  such  elements  an  excellent  method 
is  based  upon  the  study  of  specific  heats ;  while  for  the 
determination  of  the  molecular  weights  of  solid  substances 
which  can  be  dissolved  without  decomposition  a  method 
has  quite  recently  come  into  play  which  is  based  upon 


428  INORGANIC  CHEMISTRY. 

the  extent  to  which  the  compound  raises  the  boiling, 
point  or  lowers  the  freezing-point  of  its  solution.  Both 
these  methods  will  be  briefly  described  in  this  chapter. 

Next,  it  is  found  that  there  is  a  limit  to  the  law  of 
multiple  proportions.  "While,  according  to  this  law,  the 
masses  of  any  element  which  unite  with  a  given  mass  of 
another  element  bear  simple  relations  to  one  another, 
the  law  is  silent  as  to  frow  many  kinds  of  compounds  are 
possible  between  any  two  elements.  A  careful  examina- 
tion of  the  composition  of  the  compounds  of  the  ele- 
ments shows,  however,  that  there  is  a  limit  to  the  num- 
ber of  atoms  of  one  element  which  can  combine  with 
one  atom  of  another  element.  This  limit  is  determined 
by  what  is  called  the  valence  of  the  elements.  Observa- 
tions on  the  composition  of  compounds  led  to  the  hy- 
pothesis of  the  linking  of  atoms — the  linking  taking  place 
according  to  the  laws  of  valence.  The  arrangement  of 
the  atoms  in  a  molecule  is,  according  to  this,  the  consti- 
tution of  a  compound. 

Yalence,  as  we  have  seen,  is  not  a  constant  property 
of  the  atoms.  Towards  oxygen  the  elements  which  we 
have  thus  far  studied  have  the  highest  valence ;  towards 
hydrogen  the  lowest ;  and,  in  general,  towards  tjie  mem- 
bers of  the  chlorine  group  they  exhibit  an  intermediate 
valence.  The  valence  towards  hydrogen  is  in  most 
cases  constant,  while  the  valence  towards  oxygen  and 
towards  the  members  of  the  chlorine  group  varies,  in 
some  cases  between  comparatively  wide  limits,  as  between 
1  and  7  in  the  chlorine  group,  and  between  2  and  6  in  the 
sulphur  group.  Further,  the  variations  in  the  valence 
of  an  element  generally  take  place  from  odd  to  odd  or 
from  even  to  even.  In  the  case  of  chlorine  it  appears  to 
vary  from  1  to  3  to  5  to  7 ;  in  that  of  sulphur,  from  2  to 
4  to  6  ;  in  that  of  phosphorus,  from  3  to  5.  A  knowledge 
of  the  valence  of  the  elements  is  of  great  assistance  in 
dealing  with  their  compounds,  as,  knowing  their  valence, 
we  know  in  general  the  composition  of  their  principal 
compounds. 

A  comparison  of  the  atomic  weights  finally  led  to  the 
discovery  that  the  properties  of  the  elements  are  a  pen- 


CHEMICAL  ACTION.  429 

odic  function  of  these  weights.  This  is  the  great  periodic 
laiv  of  chemistry.  This  makes  a  systematic  classification 
of  the  elements  according  to  their  atomic  weights  and 
their  properties  possible,  and  is  full  of  suggestion  as  to 
the  relations  which  the  forms  of  matter  we  call  elements 
bear  to  one  another. 

Classification  of  Reactions  of  the  Elements  and  Com- 
pounds Studied. — While  there  is  undoubtedly  something 
confusing  in  the  number  of  the  compounds  and  their  reac- 
tions which  we  have  been  studying,  still,  when  these  are 
interpreted  in  the  light  of  the  atomic  theory,  of  the  law  of 
valence,  and  of  the  periodic  law,  the  study  is  much  sim- 
plified, and  those  things  which  seem  to  have  little  or  no 
connection  are  found  to  form  parts  of  a  general  system. 
In  studying  chemistry,  one  of  the  first  things  to  be  done 
is  to  learn  how  elements  and  compounds  act  upon  one 
another,  and  what  products  are  formed.  The  question 
of  composition  is  one  of  the  first  which  presents  itself, 
and  this  must  be  studied  before  other  questions  can  be 
intelligently  discussed.  What,  then,  are  the  most  promi- 
nent facts  which  we  have  learned  in  studying  the  ele- 
ments and  compounds  which  have  thus  far  been  taken 
up? 

In  the  first  place,  it  will  have  been  noticed  that,  gener- 
ally speaking,  the  compounds  which  any  element  forms 
with  oxygen  and  hydrogen  are  the  most  prominent ;  that, 
taking  the  maximum  oxygen  compound  of  an  element  as 
one  end  of  a  series,  the  other  end  is  formed  by  the  hy- 
drogen compound.  These  end-products  in  the  case  of 
chlorine,  sulphur,  phosphorus,  and  silicon  are  : 

Hydrogen  compound.  Maximum  oxygen  compound. 

HC1  C12O7 

H2S  S03(S206) 

H3P  PA 

H4Si  Si02(Si2O4) 

The  valence  towards  hydrogen  increases  while  that 
towards  oxygen  decreases  regularly  in  the  order  given. 
With  water  these  oxides  form  the  acids  HC1O4,  H2SO4, 
H.PO4,  and  H4SiO4.  Here  the  remarkable  fact  is  ob- 


430  INORGANIC  CHEMISTRY. 

served,  that  the  number  of  hydrogen  atoms  in  each  oi 
these  acids  is  the  same  as  that  in  the  hydrogen  com- 
pounds, and  the  limit  of  the  addition  of  oxygen  is  reached 
in  each  case  with  four  atoms  of  oxygen.  Further,  each 
of  the  first  three  acids  appears  to  be  related  to  so-called 
normal  acids  which  are  formed  by  union  of  the  chlorine, 
sulphur,  and  phosphorus  with  a  number  of  hydroxyl 
groups  corresponding  to  the  oxygen  valence.  These 
normal  acids  are 

Cl(OH),,    S(OH)6)    P(OH),,    Si(OH),. 

Now,  whenever  a  chlorine  compound  is  subjected  to 
oxidation  under  proper  circumstances  the  final  product 
is  perchloric  acid,  which  when  isolated  has  probably  the 
composition  represented  by  the  formula  HC1O4.  So 
when  a  sulphur  compound  is  oxidized  the  final  product 
is  sulphuric  acid,  H2SO4 ;  when  a  phosphorus  compound 
is  fully  oxidized  the  final  product  is  phosphoric  acid,, 
H3PO4 ;  and  the  final  product  of  oxidation  of  a  silicon 
compound  is  silicic  acid,  H4SiO4. 

By  reduction  of  the  above  compounds  the  final  prod- 
ucts are  the  hydrogen  compounds ;  but  before  the  limit 
of  reduction  is  reached  intermediate  products  are  formed. 
All  these  intermediate  products  are  comparatively  un- 
stable, and  tend  to  take  up  oxygen  under  ordinary  cir- 
cumstances and  to  form  the  stable  derivative  of  the 
highest  oxygen  compound.  Thus  phosphites  pass  over 
into  phosphates,  sulphites  into  sulphates,  and  chlorates 
into  perchlorates  when  heated.  These  changes  are  repre- 
sented by  the  following  equations  : 

2KC1O3  =  KC1  +  KC104  +  02 ; 
4K2S03  =  K2S  +  3K2S04 ; 
4H3PO3  =  PH3  +  3H3PO4. 

The  highest  forms  are  therefore  evidently  most  stable. 
Turning  to  the  compounds  which  the  elements  of  Families 
IY,  Y,  YI,  and  YII  form  with  the  members  of  the  chlorine 
group,  attention  has  repeatedly  been  called  to  the  fact  that 


DIRECT  COMBINATION.  431 

these  are  for  the  most  part  decomposed  by  water  with 
the  formation  of  the  corresponding  hydroxyl  compounds. 

The  elements  of  Families  IY,  Y,  VI,  and  VII  do  not 
form  compounds  with  the  members  of  the  sulphur  group, 
nor  with  those  of  the  nitrogen  group,  as  readily  as  they  do 
with  hydrogen,  with  oxygen,  and  with  the  members  of  the 
chlorine  group.  Those  elements  which  have  basic  char- 
acter, however,  like  antimony  and  bismuth,  form  very 
characteristic  compounds  with  sulphur.  The  sulphur 
compounds,  in  general,  have  a  composition  similar  to 
that  of  the  oxygen  compounds  of  the  same  elements. 

Kinds  of  Chemical  Reactions. — As  was  pointed  out  in 
the  early  part  of  this  book,  all  chemical  reactions  may  be 
classified  under  three  heads  : 

(1)  Those  which  consist  in  direct  combination  ; 

(2)  Those  which  consist  in  direct  decomposition ;  and, 

(3)  Those  which  involve  the  interaction  of  two  or  more 
elements  or  compounds  and  the  formation  of  two  or  more 
compounds.     This  is  known  as  double  decomposition  or 
metathesis. 

Direct  Combination. — We  have  had  to  deal  with  a 
number  of  examples  of  each  of  these  kinds  of  reactions. 
As  examples  of  the  first  kind  already  studied  the  fol- 
lowing may  be  mentioned  : 

The  combination  of  hydrogen  and  chlorine  to  form 
hydrochloric  acid ;  the  formation  of  ammonium  chloride 
from  ammonia  and  hydrochloric  acid ;  the  formation  of 
calcium  hydroxide  from  calcium  oxide  and  water  ;  the 
formation  of  nitrogen  peroxide  from  nitric  oxide  and 
oxygen  ;  and  the  formation  of  carbon  disulphide  from 
carbon  and  sulphur. 

As  regards  the  combination  of  hydrogen  and  chlorine, 
it  should  be  remarked  that  this  act  is  the. same  in  princi- 
ple as  that  of  metathesis.  Strictly  speaking,  it  is  not  a  case 
of  direct  combination,  as  we  understand  it.  For,  as  we 
have  seen,  according  to  the  molecular  theory,  free  chlo- 
rine and  free  hydrogen  consist  of  molecules  which  are 
made  up  of  two  atoms  each.  Therefore,  when  these  ele- 
ments are  brought  together  the  molecules  are  first  de- 
composed into  atoms  before  the  act  of  union  can  take 


432  INORGANIC  CHEMISTRY. 

place.     The  two  acts  are  represented  by  the  two  equa* 
tions  following : 

C12  +  H2  =  Cl  +  Cl  +  H  +  H ; 
d  -f  Cl  +  H  +  H  =  2HC1. 

In  the  case  also  of  the  union  of  hydrochloric  acid 
and  ammonia  it  appears  probable  that  a  serious  disar- 
rangement of  the  constituent  atoms  is  necessary  in  order 
that  the  act  of  combination  may  take  place.  According 
to  the  ammonium  theory,  ammonium  chloride  is  repre- 


sented by  the   formula 


H 
H 

H,  which  means  that  the 

H 

Cl 


atom  of  chlorine  and  four  atoms  of  hydrogen  are  in  com- 
bination with  the  atom  of  nitrogen.  But  in  order  that 
a  compound  of  this  constitution  may  be  formed  from 
ammonia  and  hydrochloric  acid,  it  is  necessary  that  the 
molecule  of  hydrochloric  acid  should  be  broken  down 
into  its  constituent  atoms.  So  that  this  case  of  apparent 
direct  combination  is,  as  far  as  we  can  judge,  in  reality 
more  complicated  than  it  appears,  and  should  be  repre- 
sented by  the  two  equations  : 

NH3  +  HC1  =  NH3  +  H  +  01; 
NH.  +  H  +  Cl  =  NH4C1. 

All  other  cases  of  apparent  direct  combination  are 
probably  of  the  same  character,  so  "that  it  is  doubtful 
whether  a  single  case  of  simple  direct  combination  is 
known. 

Direct  Decomposition. — As  examples  of  direct  decom- 
position the  following  cases  may  be  cited : 

The  decomposition  of  mercuric  oxide  by  means  of 
heat  into  mercury  and  oxygen ;  that  of  ammonium  chlo- 
ride into  ammonia  and  hydrochloric  acid  by  heat ;  that 
of  potassium  nitrate  into  potassium  nitrite  and  oxygen 


METATHESIS.  433 

by  'heat  ;  that  of  phosphorus  pentachloride  into  the 
trichloride  and  chlorine  by  heat  ;  that  of  ammonia  into 
hydrogen  and  nitrogen  by  continued  action  of  electric 
sparks  ;  and  that  of  nitrogen  iodide  by  contact  with  a 
solid  substance. 

On  close  examination  of  each  of  the  above  cases,  which 
are  fairly  typical  and  as  simple  as  any  that  could  be 
chosen,  it  will  be  seen  that  no  one  of  them  is  merely  a 
case  of  decomposition  ;  for  even  though  we  must  assume 
that  the  first  result  in  each  case  is  the  setting  free  of  the 
atoms  of  one  or  two  elements,  we  must  also  assume  that 
these  atoms  unite  again  to  form  other  molecules  either 
of  elements  or  compounds.  Thus,  when  mercuric  oxide 
is  decomposed  we  get  mercury  and  oxygen.  As  far  as 
can  be  determined,  the  mercury  atoms  do  not  unite  with 
each  other,  but  the  oxygen  atoms  do,  so  that  the  total  ac- 
tion involves  decomposition  and  afterwards  combination 
as  represented  in  the  equations 


In  the  case  of  the  pentachloride  of  phosphorus,  it  is 
probable  that  the  two  atoms  of  chlorine  are  first  given 
off  from  each  molecule  of  the  chloride,  leaving  a  molecule 
of  the  trichloride,  but  the  atoms  of  chlorine  afterwards 
unite  to  form  molecules  as  represented  thus  : 

PC15  =  PC13  +  Cl  +  Cl  ; 
PC13  +  Cl  +  Cl  =  PC13  +  C12. 

Similar  statements  hold  good  for  all  other  cases  of 
direct  decomposition. 

Metathesis.  —  This  is  the  most  common  kind  of  chemi- 
cal action,  and  indeed  from  what  has  been  said  in  regard, 
to  direct  combination  and  direct  decomposition  it  will  be 
seen  that  there  is  no  essential  difference  between  them 
and  metathesis.  Most  of  the  reactions  with  which  we 
have  had  to  deal  are  examples  of  double  decomposition 


434  INORGANIC  CHEMISTRY. 

or  metathesis,  as  :  The  formation  of  salts  by  the  action  of 
bases  upon  acids ;  the  formation  of  the  sulphides  of 
arsenic,  antimony,  and  bismuth  by  the  action  of  hydro- 
gen sulphide  upon  solutions  of  compounds  of  these  ele- 
ments ;  the  setting  free  of  hydrochloric  and  nitric  acids 
by  the  action  of  sulphuric  acid  upon  chlorides  and 
nitrates  ;  of  carbon  dioxide  and  oxides  of  nitrogen  by  the 
action  of  acids  upon  carbonates  and  nitrites  ;  and  of  am- 
monia by  treating  ammonium  salts  with  lime.  Among 
the  more  complicated  examples  which  have  been  dealt 
with  are :  The  action  of  sulphuric  acid  upon  potassium 
iodide,  giving  rise  to  the  formation  of  potassium  sul- 
phate, hydriodic  acid,  free  iodine,  sulphur  dioxide,  sul- 
phur, and  hydrogen  sulphide ;  the  action  of  chlorine 
upon  a  mixture  of  silicon  dioxide  and  charcoal ;  the  action 
of  silicon  fluoride  upon  water,  giving  rise  to  the  forma- 
tion of  silicic  acid  and  fluosilicic  acid ;  and  the  action 
of  phosphorus  pentachloride  upon  water,  forming  phos- 
phoric and  hydrochloric  acids.  As  simple  an  example 
of  this  kind  of  action  as  can  be  cited  is  that  of  the  for- 
mation of  hydrogen  and  potassium  chloride  from  potas- 
sium and  hydrochloric  acid  gas.  The  molecular  weight 
of  potassium  is  not  positively  known,  but,  assuming  its 
molecule  to  be  made  up  of  two  atoms,  the  action  must  be 
represented  in  this  way  : 

K2  +  2HC1  =  2KC1  +  H2. 

The  next  stage  of  complication  is  exhibited  in  the  re- 
action following : 

KI  +  HC1  =  KC1  +  HI. 

Examples  similar  to  the  latter,  but  somewhat  more  com- 
plicated, are  these : 

2KOH  +  H2S04  =  K2S04  +  HjO  ; 
CaCl2    +  H2S04  =  CaSO4  +  2HC1. 

The  Cause  of  Chemical  Reactions. — The  prime  cause  of 
chemical  reactions  is  something  which  we  think  of  as  an 


AN  IDEAL  CHEMICAL  REACTION.  435 

attractive  force  exerted  in  different  degrees  between  the 
different  elements.  When  any  elements  or  compounds 
are  brought  together  under  certain  conditions  the  ten- 
dency is  always  towards  the  formation  of  the  most  stable 
compounds  of  those  elements  which  can  be  formed  un- 
der the  given  conditions.  Thus,  potassium  sulphate  and 
water  are  more  stable  forms  of  combination  of  the  ele- 
ments hydrogen  and  oxygen,  and  potassium,  sulphur  and 
oxygen,  than  sulphuric  acid  and  potassium  hydroxide 
are  under  the  conditions  under  which  the  action  takes 
place.  So  also  the  system  composed  of  potassium  chlo- 
ride and  hydriodic  acid  is  more  stable  than  that  com- 
posed of  potassium  iodide  and  hydrochloric  acid  under 
the  conditions  of  the  action.  Why  the  one  system  is 
more  stable  than  the  other  we  do  not  know,  for  we  do 
not  know  what  relations  exist  between  the  atoms  in  the 
molecules.  It  is  convenient  to  think  of  that  which 
causes  the  atoms  to  unite  to  form  compounds  as  chemical 
affinity.  It  is  evident  that  this  affinity  is  more  strongly 
exerted  between  some  elements  than  between  others. 
The  affinity  of  chlorine  for  hydrogen  is,  for  example, 
much  stronger  than  that  of  chlorine  for  nitrogen  or  for 
oxygen.  Owing,  however,  to  the  complicated  character 
of  most  chemical  reactions,  it  is  extremely  difficult  to 
make  measurements  of  the  affinities  of  the  elements,  and 
but  little  progress  has  been  made  in  this  direction.  Still, 
one  of  the  great  objects  in  view  in  the  study  of  chemical 
phenomena  is  to  learn  as  much  about  chemical  affinity 
as  possible. 

An  Ideal  Chemical  Reaction. — In  every  case  in  which 
two  compounds  act  upon  each  other  to  form  two  new 
ones,  several  forces  must  be  at  work,  as  we  have  seen. 
Suppose,  for  example,  AB  and  CD  act  upon  each  other 
in  the  gaseous  condition  to  form  two  compounds  BC  and 
AD,  also  both  gaseous.  The  normal  course  of  such  a 
reaction  would  lead  to  the  formation  of  not  only  the  two 
compounds  BC  and  AD,  but  AB  and  CD  would  also  be 
present  in  the  resulting  system.  For  A  has  an  affinity 
for  B  as  well  as  for  D,  and  C  has  an  affinity  for  D  as 
well  as  for  B.  In  the  system  we  should  have  operating 


436  INORGANIC  CHEMISTRY. 

the  affinity  of  A  for  B,  and  A  for  D  ;  of  C  for  D,  and  oi 
C  for  J?.  As  these  operate  simultaneously,  equilibrium  is 
established  when  certain  quantities  of  the  four  possible 
compounds  are  formed,  the  quantities  depending  in  the 
first  instance  upon  the  relative  strengths  of  the  various 
affinities.  The  same  remarks  apply  to  the  case  in  which 
two  substances  react  in  solution  and  form  two  products 
which  are  soluble.  Here  the  action  is  not  complete  in 
any  one  direction,  but  an  equilibrium  is  established  be- 
tween the  four  possible  compounds. 

Influence  of  Mass. — The  proportions  between  the  pro- 
ducts formed  in  any  given  case  is  markedly  influenced 
by  the  relative  masses  of  the  reacting  substances.  Thus, 
sulphuric  acid  acts  upon  potassium  nitrate  when  the 
acid  is  in  excess,  forming  primary  potassium  sulphate, 
KHSO4r  and  nitric  acid.  On  the  other  hand,  if  a  large 
excess  of  nitric  acid  is  allowed  to  act  upon  primary 
potassium  sulphate,  sulphuric  acid  and  potassium  nitrate 
are  produced.  Much  attention  has  been  given  to  the 
study  of  mass  action,  and  the  result  is  to  show  that  in 
reactions  generally  this  kind  of  action  comes  prom- 
inently into  play.  The  law  has  been  established  that 
chemical  action  is  proportional  to  the  product  of  the  active 
masses  of  the  substances  taking  part  in  the  change.  It  would 
appear  from  this  that  the  decomposition  of  two  com- 
pounds to  form  two  new  ones  would  not  be  complete,  if 
the  conditions  are  such  that  the  two  new  compounds  can 
act  upon  each  other.  If  a  large  excess  of  one  of  the  re- 
acting compounds  is  taken,  however,  the  reaction  may  be 
made  approximately  complete  by  reason  of  the  mass 
action. 

Reactions  May  be  Complete  if  one  of  the  Products 
Formed  is  Insoluble  or  Volatile. — When  two  substances 
which  by  interaction  can  form  an  insoluble  product  are 
brought  together,  the  reaction  generally  takes  place  and 
is  complete.  When  the  substances  are  brought  together 
we  may  imagine  that,  owing  to  interaction,  a  small  quan- 
tity of  the  insoluble  compound  is  formed  at  once.  If 
this  product  were  soluble,  the  action  would  stop  before 
it  is  complete,  because  this  new  product  would  itself 


WHEN  REACTIONS  MAT  BE  COMPLETE.  437 

exert  its  action  upon  the  system.  Being  insoluble,  how- 
ever, it  is  removed  from  the  sphere  of  action,  and  the 
same  reaction  which  caused  the  formation  of  the  first 
particles  of  it  can  now  be  repeated,  and  so  on,  until  the 
reaction  is  complete.  This  is  illustrated  in  the  action  of 
sulphuric  acid  upon  barium  chloride  in  solution.  The 
two  substances  react  as  represented  in  this  equation : 

BaCl2  +  H3S04  +  Aq  =  BaSO4  +  HC1  +  Aq. 

The  symbol  Aq  is  simply  intended  to  indicate  that  the 
reaction  takes  place  in  solution.  If  barium  sulphate  were 
soluble,  all  four  substances — barium  chloride,  sulphuric 
acid,  barium  sulphate,  and  hydrochloric  acid — would  be 
present  in  the  solution  after  the  establishment  of  equi- 
librium. But,  being  insoluble,  it  is  removed,  and  new 
quantities  are  formed  as  long  as  the  substances  necessary 
for  its  formation  are  present  in  the  solution  ;  that  is,  until 
either  all  the  barium  chloride  is  decomposed  or  all  the 
sulphuric  acid  is  removed.  Reactions  involving  the  for- 
mation of  insoluble  compounds  or  precipitates  are  among 
the  most  common  with  which  we  have  to  deal,  particu- 
larly in  the  various  operations  of  analytical  chemistry. 

Again,  when  two  substances  which  can  form  a  volatile- 
product  are  brought  together  the  reaction  generally  takes 
place  and  is  complete.  The  reason  why  a  reaction  of 
this  kind  is  complete  is  the  same  as  that  given  in  the 
case  of  the  formation  of  an  insoluble  compound.  Each 
successive  portion  of  the  volatile  product  formed  is  re- 
moved, and  the  reaction  which  gave  rise  to  it  proceeds  as 
long  as  the  necessary  substances  are  present.  This  kind 
of  action  has  been  repeatedly  illustrated.  It  is  that,  for 
example,  which  is  seen  in  the  liberation  of  hydrochloric 
acid  from  a  chloride  by  the  action  of  sulphuric  acid ;  of 
carbon  dioxide  by  the  action  of  an  acid  upon  a  carbonate ; 
and  of  ammonia  by  the  action  of  lime  upon  ammonium 
chloride. 

An  interesting  example  of  the  combined  influence  of 
mass  and  the  volatility  of  the  product  is  seen  in  the  action 
of  heated  iron  upon  an  excess  of  steam,  and  of  the  oxide 


438  INORGANIC  CHEMISTRY. 

of  iron  upon  an  excess  of  hydrogen.  When  steam  is 
passed  over  heated  iron,  action  takes  place  thus  : 

4HaO  +  3Fe  =  Fe3O4  +  4H2. 

Hydrogen  is  liberated  and  the  oxide  of  iron  formed. 
When,  however,  hydrogen  is  passed  over  heated  oxide 
of  iron  the  reverse  reaction  takes  place  : 

Fe3O4  +  4H2  =  3Fe  +  4H2O. 

Owing  to  the  excess  of  steam  always  present  in  the  first 
reaction,  hydrogen  is  constantly  formed  and  constantly 
being  removed.  Undoubtedly  the  hydrogen  formed  acts 
to  some  extent  upon  the  oxide,  but  the  other  reaction 
always  takes  place  to  a  greater  extent.  The  opposite  is 
true  when  the  oxide  is  heated  in  an  excess  of  hydrogen. 
The  principal  reaction  which  takes  place  in  this  case  is 
that  of  the  hydrogen  upon  the  oxide  of  iron,  and  the 
steam  is  carried  out  of  the  field  almost  as  soon  as  formed, 
so  that  the  reduction  of  the  oxide  of  iron  continues. 

Thermochemical  Study  of  Affinity . — If  a  mass  of  hydro- 
gen and  a  mass  of  chlorine  consisted  of  isolated  atoms  at 
rest,  and,  after  combination,  the  molecules  as  well  as 
their  constituent  atoms  were  at  rest,  then  the  heat  evolved 
in  the  act  of  combination  would  be  the  result  of  the  trans- 
formation of  the  potential  energy  of  the  atoms  into  kinetic 
energy,  and  it  would  be  a  measure  of  the  affinity  exerted 
between  the  atoms.  But  none  of  these  conditions  can 
be  assumed  with  any  confidence,  and  most  of  them  un- 
doubtedly do  not  exist.  We  have  abundant  evidence  to 
show  that  the  mass  of  hydrogen  and  that  of  chlorine  do 
not  consist  of  isolated  atoms.  Taking,  then,  the  reaction 
between  hydrogen  and  chlorine,  it  is  clear,  as  lias  already 
been  explained,  that  it  is  not  simply  a  combination  of 
atoms,  but  that  the  act  of  combination  between  the  atoms 
must  be  preceded  by  the  decomposition  of  the  molecules 
of  hydrogen  and  those  of  chlorine.  The  heat  which  is 
evolved  in  the  reaction  is  therefore  not  simply  the  result 
of  the  combination  of  hydrogen  and  chlorine,  but  it  is 


VALUE  OF  THERMOCHEMICAL  MEASUREMENTS.    439 

this  heat  less  that  which  is  required  to  decompose  the 
molecules  of  hydrogen  and  those  of  chlorine  into  atoms. 
The  heat  measured  is  the  difference  between  two  quan- 
tities ;  and  we  have  no  means  of  estimating  the  value  of 
these  quantities.  This  is  true  of  every  chemical  reaction. 
The  heat  evolved  or  absorbed  in  the  reaction  is  the  dif- 
ference between  two  or  more  quantities,  and  it  is  not 
therefore  a  measure  of  affinity. 

Nevertheless,  some  knowledge  regarding  the  relations 
which  the  affinities  of  elements  bear  to  one  another  can 
be  gained  by  a  study  of  the  heat  evolved  in  their  re- 
actions. Thus,  the  following  results  have  been  obtained 
in  the  study  of  chlorine,  bromine,  and  iodine  in  their  re- 
lations to  hydrogen : 

[H2,  C1J  =  2[H,  Cl]  -  [H,  H]  -  [01,  Cl]  =  44,000  c. 
[H2,  BrJ  =  2[H,  Br]  -  [H,  H]  -  [Br,  Br]  =  16,880  c. 
[H2,I3]  =2[H,I]  -  [H,  H]  -  [I,  I]  =  12,072  c. 

The  meaning  of  these  three  equations  will  appear  from 
an  interpretation  of  the  first.  This  means  that  when  a 
molecule  of  hydrogen  acts  upon  a  molecule  of  chlorine 
to  form  two  molecules  of  hydrochloric  acid  gas  44,000  c. 
of  heat  are  evolved ;  and  this  quantity  is  the  difference 
between  that  which  is  evolved  in  the  combination  of  two 
atoms  of  hydrogen  with  two  atoms  of  chlorine,  and  that 
which  is  absorbed  in  the  decomposition  of  one  molecule 
of  hydrogen  into  two  atoms,  and  in  the  decomposition  of 
one  molecule  of  chlorine  into  two  atoms.  The  figures 
thus  obtained  are  not  proportional  to  the  affinities  of 
chlorine,  bromine,  and  iodine  for  hydrogen,  but  never- 
theless the  affinities  in  all  probability  vary  in  the  same 
order. 

The  difficulties  are  much  increased  in  more  complicated 
cases,  and  it  will  therefore  be  seen  that  it  is  impossible 
to  measure  the  affinity  between  the  atoms  by  means  of 
the  heat  evolved  in  reactions. 

Value  of  Thermochemical  Measurements. —  Although 
the  affinities  of  the  elements  for  one  another  cannot  be 
directly  estimated  by  means  of  thermochemical  measure- 


440  INORGANIC  CHEMISTR  Y. 

ments,  nevertheless  these  measurements  are  valuable,  as 
they  show  a  direct  relation  between  the  quantity  of  heat 
evolved  and  the  character  of  the  reaction  which  takes 
place  in  any  given  case.  In  the  case  above  cited,  for 
example,  it  is  seen  that  the  heat  of  formation  of  hydro- 
chloric acid  is  greater  than  that  of  hydrobromic  acid, 
and  that  of  hydrobromic  acid  is  in  turn  greater  than 
that  of  hydriodic  acid.  Now,  on  page  94,  it  was  stated 
that  in  general  that  exothermic  reaction  takes  place 
which  is  accompanied  by  the  greatest  evolution  of  heat.* 
Accordingly,  in  a  case  in  which  both  hydrochloric  and 
hydrobromic  acid  could  be  produced  the  former  would 
certainly  be  produced  in  larger  quantity. 

Heat  of  Neutralization— Avidity  of  Acids. — Among  the 
measurements  which  have  proved  of  value  in  connection 
with  the  study  of  the  general  problem  of  affinity,  are 
those  furnished  by  the  heat  of  neutralization  of  acids 
and  bases.  The  general  method  of  work  consisted  in 
determining  the  heat  evolved  when  equivalent  quantities 
of  different  acids  are  neutralized  by  the  same  base  and 
equivalent  quantities  of  different  bases  are  neutralized  by 
the  same  acid.  Knowing  the  heat  evolved  in  the  reactions 
between  the  various  acids  and  bases,  it  .is  possible  to  de- 
termine what  takes  place  when  acids  act  upon  salts  in 
which  decomposition  is  not  evident,  either  from  the  for- 
mation of  a  precipitate  or  the  evolution  of  a  gas.  Thus, 
when  two  molecules  of  nitric  acid  act  upon  one  of  sodium 
sulphate  in  solution,  several  changes  are  possible,  as 
represented  in  the  equations 

(1)  Na2SO4  +    HNO3  =  NaHS04  +    NaNO3 ; 

(2)  Na2S04  +  2HNO3  =  H2SO4      +  2NaNO3 ; 

(3)  2Na2SO4  +  4HNO3  =  Na2SO4    +  2NaNO3  +  H2SO4 

+  2HNO3. 

As  all  the  substances  involved  in  these  reactions  are 
soluble  in  water,  and  the  reactions  are  studied  in  water 
solution,  it  is  clear  that  by  ordinary  methods  it  would 
be  impossible  to  tell  which  of  them  takes  place.  By 
measuring  the  heat  evolved,  however,  it  has  been  shown 

*  This  statement  is  known  as  Berthelot's  Law  of  Maximum  Work. 


AVIDITY  OF  ACIDS.  441 

that  in  this  and  in  all  similar  cases  the  base  is  divided 
between  the  two  acids,  and  generally  more  goes  to  one 
acid  than  to  the  other.  Further,  it  is  possible  to  meas- 
ure the  division  of  the  base  between  the  acids,  and  in 
this  way  measurements  of  the  relative  strengths  of  acids 
are  obtained.  The  figures  representing  the  strengths  of 
the  acids  measured  in  this  way  are  called  the  avidities  of 
the  acids.  In  the  case  taken  above  as  an  illustration,  it 
was  found  that  in  dilute  aqueous  solution  two-thirds  of 
the  sodium  goes  to  the  nitric  acid  and  one-third  to  the 
sulphuric  acid.  Therefore,  it  appears  that  the  avidity 
of  nitric  acid  is  twice  as  great  as  that  of  sulphuric 
acid.  Of  all  acids  investigated,  nitric  and  hydrochloric 
acids  were  found  to  have  the  greatest  avidity.  Calling 
this  100,  the  avidities  of  some  other  acids  determined  by 
this  method  are  as  given  in  the  following  table  : 

Acids.  Avidity. 

Nitric  acid, 100 

Hydrochloric  acid, 100 

Hydrobromic  acid,       ......  89 

Hydriodic  acid, •  V  79 

Sulphuric  acid, ,  49 

Selenic  acid, 45 

Hydrofluoric  acid, 5 

Boric  acid,       .     .     «     .     .     ,     .     .     „  1 

Silicic  acid,     ..........  0 

Hydrocyanic  acid, 0 

The  figures  given  refer  to  equivalent  quantities  of  the 
acids,  i.e.,  quantities  which  can  be  neutralized  by  equal 
quantities  of  a  base.  Thus,  1  molecule  of  nitric  acid, 
HNO3,  is  neutralized  by  1  molecule  of  sodium  hydroxide, 
NaOH  ;  but  only  £  molecule  of  sulphuric  acid  is  neutral- 
ized by  1  molecule  of  sodium  hydroxide,  and  only  % 
molecule  of  orthophosphoric  acid  would  be  neutralized 
by  the  same  quantity  of  base.  Therefore,  we  say  that 
1  molecule  of  nitric  acid  is  equivalent  to  £  molecule  of 
sulphuric  acid,  and  to  -J-  molecule  of  orthophosphoric 
acid. 


442  INORGANIC  CHEMISTRY. 

It  is  impossible  at  present  to  give  an  exact  interpretation 
of  the  results  above  recorded,  but  it  appears  that  the 
figures  given  represent  the  numerical  relations  between 
some  common  property  possessed  by  acids,  a  property 
which  we  have  vaguely  in  mind  when  we  speak  of  the 
strength  of  acids.  This  appears  more  clearly  when  acids 
and  bases  are  studied  in  other  ways. 

Other  Methods  for  Determining  the  Avidity  of  Acids. 
— Besides  the  thermochemical  method  of  studying  the 
action  of  acids  on  bases,  several  other  methods  have  been 
devised.  Among  these  are  the  volume-chemical  method, 
the  optical  method,  the  action  of  acids  on  insoluble  salts, 
and  the  electrical  method.  The  object  in  view  is  in  all 
cases  practically  the  same — to  compare  the  influence 
exerted  by  different  acids  under  the  same  circumstances, 
and  thus  to  measure  their  avidity  or,  as  this  has  also 
been  called,  their  specific  coefficient  of  affinity. 

(1)  The  volume-chemical   method  depends  upon  the 
fact  that  chemical  processes  which  take  place  in  homo- 
geneous  liquids    generally    cause   changes    in   volume. 
"  Thus,  the  specific  gravity  of  a   normal   caustic  soda 
solution  was  found  to  be  1.04051,  that  of  an  equivalent 
solution  of  sulphuric  acid  1.0297,  that  of  an  equivalent 
of  nitric   acid   1.03089.      When  equal  volumes  of  soda 
solution  were  mixed  with  each  of  the  acids,  the  specific 
gravity  of  the  sodium  sulphate  solution  was  1.02959,  and 
that  of  the  nitrate  solution  1.02633.     Finally,  when  to 
the  solution  of  sodium  sulphate  (2  vols.)  one  equivalent 
(1  vol.)  of  nitric  acid  was  added,  the  specific  gravity  be- 
came 1.02781."     By  means  of  these  figures  it  is  possible 
to  determine  to  what  extent  the  nitric  acid  acts  upon  the 
sulphate,  and  thus  to  draw  conclusions  regarding  the 
distribution  of  the  base  between  the  acids.     The  results 
reached  by  this  method  agree  in   general  with   those 
reached  by  the  thermochemical  method. 

(2)  In  the  optical  method  the  coefficient  of  refraction 
of  various  solutions  is  determined,  and  also  the  changes 
in  the  coefficient  of  refraction  produced  by  mixing  these 
solutions  in  certain  ways,  and  thus  it  is  possible  to  draw 


DISSOCIATION.  443 

conclusions  in  regard  to  the  character  of  reactions  which 
take  place  in  solutions. 

(3)  An  illustration  of  the  method  involving  the  action 
of  acids  on  insoluble  salts  will  make  the  method  clear. 
A  weighed  quantity  of  calcium  oxalate  is  treated  with 
equivalent  quantities  of  different  acids  in  dilute  solutions, 
and  the  quantity  of  the  salt  dissolved  in  a  given  time 
then  determined.      From  the   result   it   is   possible   to 
calculate  the  specific  coefficients  of  afiinity  of  the  acids. 

(4)  The  simplest  method  of  all  is  the  electrical.     This 
consists  in  determining  the  conducting  power  of  different 
substances  at  the  same  dilutions.     In  this  way  figures 
are*  obtained  which  bear  to  one  another  the  same  rela- 
tions as  those  expressing  the  coefficients  of  affinity. 

It  is  impossible  to  go  into  details  in  regard  to  these 
methods  here,  and  it  need  only  be  said  that  when  acids 
and  bases  are  compared  by  the  above  methods,  they  are 
found  to  differ  markedly  from  one  another,  and  the  order 
in  which  they  are  arranged  by  the  results  of  the  different 
methods  is  always  essentially  the  same. 

Study  of  Chemical  Decompositions. — As  we  have  seen, 
practically  every  case  of  chemical  combination  with  which 
we  have  to  deal  is  associated  with  the  decomposition  of 
molecules,  so  that  it  is  impossible  perfectly  to  sepa- 
rate the  two  acts  of  combination  and  decomposition. 
Nevertheless  there  are  some  comparatively  simple  cases 
of  decomposition  which  have  been  studied  with  special 
care,  and  results  of  much  importance  have  been  obtained. 
The  most  interesting  are  those  cases  of  decomposition 
which  are  included  under  the  heads  of  dissociation  and 
electrolysis.  While  many  chemical  decompositions  are 
brought  about  by  concussion — that  is,  by  mechanical  dis- 
turbance of  the  mass — the  very  instability  of  the  com- 
pounds which  makes  these  decompositions  possible  pre- 
vents any  very  profitable  study  of  the  phenomena. 

Dissociation. — Attention  has  been  called  to  the  fact  that 
many  compounds,  when  heated  to  sufficiently  high  tem- 
peratures, are  decomposed.  Thus,  water  is  partly  de- 
composed into  hydrogen  and  oxygen  when  heated  to  1000°; 


444  INORGANIC  CHEMISTRY. 

ammonium  chloride  is  decomposed  into  ammonia  and  hy- 
drochloric acid ;  phosphorus  pentachloride,  into  the  tri- 
chloride and  chlorine  ;  nitrogen  peroxide  of  the  formula 
N2O4,  into  the  simpler  compound  of  the  formula  NO2,  etc. 
Careful  study  of  any  one  of  these  cases  shows  the  follow- 
ing facts  :  (1)  That  the  decomposition  takes  place  gradu- 
ally ;  (2)  that  the  extent  of  the  decomposition  depends 
upon  the  temperature  and  pressure,  and  for  the  same 
compound  is  always  the  same  for  the  same  temperature 
and  pressure  ;  (3)  that  if  the  full  amount  of  decomposition 
possible  at  a  certain  temperature  is  effected,  and  the  tem- 
perature then  lowered,  the  constituents  will  recombine  to 
some  extent  until  equilibrium  at  the  lower  temperature  is 
established. 

In  a  case  of  dissociation  by  heat,  then,  the  decomposi- 
tion is  carried  farther  and  farther  as  the  temperature  is 
raised  higher  and  higher,  and  it  is  finally  complete.  On 
lowering  the  temperature  again,  more  and  more  of  the 
compound  is  formed  by  the  recombination  of  the  constit- 
ents  until,  when  the  lower  temperature  is  again  reached, 
there  is  no  decomposition. 

The  explanation  of  the  phenomenon  of  dissociation  is 
found  in  the  kinetic  theory  of  gases.  According  to  this 
theory,  the  molecules  of  a  gas  at  a  given  temperature  are 
moving  with  different  velocities,  though  the  average 
velocity  of  all  the  molecules  is  the  same  at  the  same 
temperature.  Now,  it  is  highly  probable  that  the  motion 
of  the  atoms  within  the  molecules  partakes  of  that  of 
the  molecules  themselves,  so  that  the  motion  of  the 
atoms  in  the  molecules  with  the  greatest  velocity  is 
probably  the  greatest,  and,  in  these,  decomposition 
will  take  place  first.  When  a  compound  gas  is  heated, 
we  can  easily  conceive  that  even  at  a  comparatively  low 
temperature  the  motion  of  some  of  the  molecules  will 
be  sufficient  to  cause  their  decomposition,  and,  as  the 
average  motion  of  all  the  molecules  is  constant  for  a 
given  temperature,  the  amount  of  decomposition  will 
be  constant  for  that  temperature.  As  the  molecules  are, 
however,  moving  in  every  direction  and  constantly  col- 
liding, a  molecule  which  is  decomposed  at  one  instant 


ELECTROLYSIS.  445 

may  be  re-formed  at  the  next,  and  one  that  is  not  decom- 
posed may  acquire  motion  enough  to  cause  its  decompo- 
sition. Though,  as  is  believed,  these  changes  are  con- 
stantly taking  place  at  every  temperature,  still,  as  has 
been  said,  the  number  of  molecules  which  will  be  decom- 
posed in  a  given  mass  at  a  given  temperature  and  pres- 
sure will  always  be  the  same.  The  higher  the  tem- 
perature, then,  the  greater  the  number  of  molecules  in 
the  conditions  which  cause  decomposition,  and  the 
smaller  the  number  of  those  in  the  conditions  favorable 
to  formation.  At  each  temperature  and  pressure  an 
equilibrium  is  established,  the  number  of  molecules  de- 
composed being  equal  to  the  number  formed.  It  is 
obvious  that,  if  one  of  the  products  of  decomposition  is 
removed,  the  conditions  are  entirely  changed.  Then  the 
possibility  of  recombination  will  not  exist,  and  total 
decomposition  can  be  effected  at  a  lower  temperature 
than  that  required  for  total  decomposition  in  the  process 
of  dissociation  proper. 

Electrolysis. — Some  chemical  compounds  in  solution 
in  water  conduct  electricity,  and  at  the  same  time  they 
undergo  decomposition.  Thus,  hydrochloric  acid  in  solu- 
tion in  water  conducts  electricity  and  the  compound  is 
decomposed  into  its  constituents  hydrogen  and  chlorine, 
the  hydrogen  appearing  at  the  negative  and  the  chlorine 
at  the  positive  pole.  Compounds  that  act  in  this  way 
are  called  electrolytes.  When  a  current  of  electricity 
acts  upon  solutions  of  different  salts,  equivalent  quanti- 
ties of  the  metals  are  deposited  by  the  same  current  in 
the  same  time.  This  is  Faraday's  Law.  Thus  if  the  same 
current  were  passed  simultaneously  through  solutions 
of  silver  nitrate,  AgNO3 ,  mercuric  nitrate,  Hg(NO3)a , 
cupric  sulphate,  CuSO4,  and  ferric  chloride  FeCls,  it 
would  be  found  that  for  every  107.11  parts  by  weight 
of  silver  deposited  there  would  be  99.2  parts  by  weight 
of  mercury  deposited,  31.56  of  copper,  and  18.53  of 
iron.  These  are  equivalent  quantities  of  these  metals 
—  quantities  that  take  the  place  of  one  part  by 
weight  of  hydrogen — and  are  to  be  distinguished  from 
atomic  quantities.  Those  elements  which  appear  at 


446  INORGANIC  CHEMISTRY. 

the  negative  pole  are  called  electro-positive,  and  those  which 
appear  at  the  positive  pole  are  called  electro-negative. 
Those  elements  which  we  call  acid-forming  are  electro- 
negative, while  hydrogen  and  the  base-forming  elements 
are  electro-positive.  The  electrolysis  of  chemical  com- 
pounds is  not  generally  a  simple  decomposition  into  two 
constituents.  Thus,  when  copper  sulphate,  CuSO4,  is 
decomposed,  the  copper  is  deposited  at  the  negative 
pole ;  but  no  such  compound  as  SO4  appears  at  the  posi- 
tive pole.  This,  if  formed,  perhaps  breaks  down  into 
oxygen  and  sulphur  trioxide,  and  the  latter  with  water 
would  then  naturally  form  sulphuric  acid.  Both  oxygen 
and  sulphuric  acid  as  a  matter  of  fact  appear  at  the  posi- 
tive pole.  The  changes  involved  may  be  represented  thus,: 

CuSO4=Cu+S04;     S04=S03+0;    SO3+HaO=H2SO4. 

Electrolytic  Dissociation. — It  has  been  known  for  a 
long  time  that  a  very  weak  electric  current  acting  upon  a 
solution  of  an  electrolyte  is  sufficient  to  cause  the  ions 
to  appear  at  the  poles.  This  fact  is  inexplicable  if 
it  is  assumed  that  the  current  is  the  cause  of  the  de- 
composition of  the  electrolyte.  This  and  some  other 
facts  which  will  be  referred  to  farther  on  make  it  prob- 
able that  electrolytes  are  at  least  to  some  extent  decom- 
posed into  their  constituent  ions  when  they  are  dissolved  ; 
that  these  ions  charged  with  electricity  transfer  their 
charges  in  the  solution  and  thus  conduct  the  current ;  and 
that  when  an  ion  charged  with  negative  electricity  reaches 
the  positive  pole  its  electricity  is  discharged,  and  the 
ion  then  ceases  to  be  an  ion  and  becomes  an  element  in 
the  free  state  or  some  compound  which  appears  either  as 
such  or  in  the  form  of  other  products.  According  to 
this  conception,  the  act  of  solution  of  an  electrolyte,  in 
water  at  least,  involves  partial  breaking  down  or  dis- 
sociation of  the  compound  into  its  ions.  The  extent  of 
this  breaking  down  is  determined  primarily  by  the  con- 
centration of  the  solution — the  greater  the  dilution  the 
greater  the  dissociation.  At  infinite  dilution  there  is 
complete  dissociation.  A  water  solution  of  hydrochloric 


ELECTROLYTIC  DISSOCIATION.  447 

acid  containing  36.18  grams  of  the  acid  in  1000  liters 
has  been  shown  to  be  completely  dissociated,  or  it  is  to 
be  regarded  as  containing  ions  of  hydrogen  and  of 
chlorine.  These  and  all  other  ions  are  carefully  to  be  dis- 
tinguished from  the  atoms  or  definite  compounds.  An 
ion  always  carries  with  it  a  certain  charge  of  electricity. 
When  this  is  discharged  the  ion  becomes  either  an  ele- 
ment or  a  compound  in  the  free  state.  When  a  solution 
of  one  electrolyte  acts  upon  a  solution  of  another  the 
reaction  observed  is  probably  due  to  the  interaction  of 
the  ions,  and  it  is  further  probable  that,  so  far  as  the 
compounds  are  present  in  the  undissociated  condition, 
they  do  not  act  upon  each  other.  If  this  view  is  correct 
the  reactions  most  familiar  to  us  are  reactions  of  ions, 
and  not  of  elements  or  compounds.  When,  for  ex- 
ample, an  acid  acts  upon  a  base  in  solution  it  appears 
that,  so  far  as  they  react,  they  are  in  dissociated  condi- 
tion. Thus  hydrochloric  acid  and  sodium  hydroxide  are 
to  be  regarded  as  acting  as  represented  in  the  following 
equation : 

H  +  01  +  Na  +  OH  =  Na  +  Cl  +  H2O. 

The  act  consists  in  the  union  of  the  hydroxyl  ion  of 
the  base  with  the  hydrogen  ion  of  the  acid  to  form  water, 
the  sodium  and  chlorine  ions  remaining  as  ions  as  in 
a  dilute  solution  of  sodium  chloride.  In  the  case  of 
nitric  acid  and  potassium  hydroxide  the  following  equa- 
tion represents  the  reaction  at  infinite  dilution : 

H  +  N03  +  K  +  OH  =  K  +  NO,  +  H,O. 

And  so  also  whenever  the  act  of  neutralization  takes 
place  there  is  simply  a  union  of  hydrogen  ions  with 
hydroxyl  ions  to  form  water. 

This  conception  finds  strong  confirmation  in  the  fact 
that  the  heat  evolved  in  neutralizing  equivalent  quantities 
of  all  acids  at  infinite  dilution  is  always  the  same — a  fact 
that  it  is  difficult  to  explain  if  it  is  assumed  that  in  the 
act  of  neutralization  a  salt  is  formed  in  the  solution. 

In  the  following  table  the  heats  of  neutralization  of 
a  few  acids  and  bases  are  given : 


448  INORGANIC1  CHEMISTRY. 

Acid  and  Base.  Heat  of 

Neutral. 

Hydrochloric  acid  and  sodium  hydrox.,  .  .  13,700 

"  "  lithium  "  .  .  13,700 

"  "  "  potassium  "  .  .  13,700 

"  "  "  barium  "  .  .  13,800 

"  "  calcium  "  .  .  13,900 

Hydrobromic    "  "  sodium  "  '  .  .  13,700 

Nitric  "  "          "  "  .  .  13,700 

lodic  "  "  "  "  .  .  13,700 

As  will  be  seen,  the  heat  of  neutralization  is  the  same 
no  matter  what  the  base  or  what  the  acid  may  be,  and 
as  has  been  pointed  out  this  fact  is  easily  understood,  if 
the  act  of  neutralization  consists  in  the  union  of  a  hy- 
droxyl  ion  with  a  hydrogen  ion  to  form  water. 

If  the  strength  of  an  acid  is  determined  by  the  extent 
to  which  it  is  dissociated,  then  of  course  those  acids  that 
are  most  readily  dissociated  are  the  strongest.  By  every 
method  available  hydrochloric  and  nitric  acids  are  found 
to  be  the  most  readily  dissociated  and  they  are  the  strong- 
est acids.  The  electrical  method  for  the  determination 
of  the  strength  of  acids  is  based  upon  this  theory.  The 
applicability  of  this  theory  to  the  explanation  of  the  most 
common  reactions  that  take  place  in  solution  is  at  present 
attracting  the  attention  of  chemists,  and  promises  to  be 
of  great  service  to  chemistry.  Those  reactions  which  are 
made  use  of  for  the  purpose  of  detecting  the  presence  of 
the  various  elements  appear  in  fact,  as  has  already  been 
stated,  to  be  reactions  of  the  ions,  and  when  these  ions 
are  not  present  the  reactions  are  not  observed.  For  ex- 
ample, when  silver  nitrate  in  solution  is  added  to  sodium 
chloride  in  solution  a  precipitate  of  silver  chloride  is 
formed,  the  reaction  taking  place  as  represented  in  this 
equation  : 

AgNO3  +  ^aCl  =  AgCl  +  NaNO,. 

This  reaction  seems,  however,  to  be  due  to  the  fact  that 
ions  of  silver  and  of  chlorine  are  present.  These  com- 
ing together  form  the  insoluble  molecules  silver  chloride 
which  is  then  precipitated.  There  are  many  compounds 
that  contain  chlorine  and  yet  do  not  give  a  precipitate 


SPECIFIC  HEAT  AND  ATOMIC  WEIGHTS.  449 

of  silver  chloride  when  treated  with  silver  nitrate.  It  is 
believed  that  in  these  cases  the  compound  is  not  ionised^ 
by  the  solvent  in  such  a  way  as  to  yield  ions  of  chlorine, 
and  that  therefore  there  are  no  chlorine  ions  present. 
An  example  of  a  compound  that  contains  chlorine  and 
yet  does  not  ordinarily  give  the  reactions  of  this  element 
is  potassium  chlorate,  KC1O3.  A  solution  of  this  salt 
does  not  give  a  precipitate  with  silver  nitrate.  It  is  pro- 
bable that  the  reason  of  this  is  that  the  salt  gives  the 
ions  K  and  C1O3 ,  the  latter  acting  quite  differently  from 
the  ion  01,  as  we  should  naturally  expect. 

Relations  between  Specific  Heat  and  Atomic  Weights.— 
The  fact  that  there  is  a  method  for  the  determination  of 
atomic  weights  founded  upon  the  relations  existing  be- 
tween these  weights,  and  the  specific  heat  of  the  ele- 
ments, has  been  mentioned.  It  has  been  found  that,  when 
equal  weights  of  different  elements  are  exposed  to  exactly 
the  same  source  of  heat,  they  require  different  lengths  of 
time  to  become  heated  to  the  same  temperature.  Given 
exactly  the  same  heating  power,  it  requires  32  times  as 
long  to  raise  the  temperature  of  a  pound  of  water  10,  20, 
or  30  degrees  as  it  does  to  raise  the  temperature  of  a 
pound  of  mercury  the  same  number  of  degrees  ;  or  it  takes 
32  times  as  much  heat  to  raise  a  pound  of  water  10,  20, 
or  30  degrees  as  it  does  to  raise  a  pound  of  mercury 
the  same  number  of  degrees.  Starting  at  the  same  tem- 
perature the  quantity  of  heat  required  to  raise  the  tem- 
perature of  a  certain  weight  of  a  substance  one  degree, 
as  compared  with  the  quantity  of  heat  required  to  raise 
the  temperature  of  the  same  weight  of  water  one  degree, 
is  called  the  specific  heat  of  the  substance.  Thus,  from 
what  was  said  above,  the  specific  heat  of  mercury  is  -£%, 
or,  in  decimals,  0.03125.  In  a  similar  way  it  can  be  shown 
that  the  specific  heat  of  gold  is  0.03244 ;  of  zinc,  0.0955  ; 
of  silver,  0.057 ;  of  copper,  0.0952. 

Now,  when  solid  elements  are  examined  with  reference 
to  their  specific  heats,  a  very  simple  relation  is  found  to 
exist  between  the  numbers  expressing  the  specific  heats 
and  the  atomic  weights.  This  relation  will  be  made 
clear  by  a  consideration  of  a  few  cases : 


450  INORGANIC  CHEMISTRY. 

Element.  Specific  Heat.  Atomic  Weight. 

Silver, 0.0570  107.11 

Zinc,.   :.>-'••  ..;-;Y.-y.y    0.0955  64.91 

Cadimum,    ....     0.0567  111.10 

Copper, 0.0952  63.12 

Tin, 0.0562  118.15 

An  examination  of  this  table  will  show  that  the  atomic 
weights  are  inversely  proportional  to  the  specific  heats. 
We  have 


107.11  :  64.91 
111.10:  63.12 
107.11  :  118.15 


0.0955  :  0.0570 ; 
0.0952  :  0.0567 ; 
0.0562  :  0.0570;  etc. 


These  proportions  are  only  approximately  correct ;  but 
it  must  be  remembered  that  the  means  for  the  determi- 
nation of  atomic  weights  and  specific  heats  are  not  per- 
fect, and  in  both  sets  of  figures  there  are  undoubtedly 
small  errors.  Hence  such  slight  variations  from  abso- 
lute agreement  in  these  proportions  should  occasion  no 
surprise.  The  agreement  is  sufficiently  close  to  indicate 
a  close  connection  between  the  two  sets  of  figures.  This 
connection  may  be  stated  in  another  way :  The  product 
of  the  atomic  weight  by  the  specific  heat  is  a  constant.  Thus, 
in  the  above  cases  : 

107.11  X  0.057    =6.11; 

64.91  X  0.0955  =  6.20 ; 
111.10  X  0.0567  =  6.30; 

63.12  X  0.0952  =  6.01 ; 
118.15  X  0.0562  =  6.64. 

From  the  above  it  appears  that  the  quantity  of  heat 
necessary  to  raise  masses  of  the  elements  proportional 
to  their  atomic  weights  the  same  number  of  degrees  is 
the  same  in  all  cases.  Suppose  two  elements  to  have 
the  atomic  weights  2  and  4.  Their  specific  heats  would 
be  to  each  other  as  2  to  1.  That  is  to  say,  it  would 
require  twice  as  much  heat  to  raise  the  temperature  of  a 
given  mass  of  the  element  with  the  atomic  weight  2  a 


SPECIFIC  HEAT  AND  ATOMIC  WEIGHTS.          451 

certain  number  of  degrees,  as  it  would  require  to  raise 
the  temperature  of  the  same  mass  of  the  element  with 
the  atomic  weight  4  the  same  number  of  degrees.  But 
to  raise  the  temperature  of  masses  of  these  two  elements 
proportional  to  their  atomic  weights  would  require  the 
same  quantity  of  heat.  This  fact  may  be  stated  thus  : 

The  atoms  of  all  elements  have  the  same  capacity  for  heat. 
This  is  only  another  way  of  stating  that,  to  raise  the 
temperature  of  an  atom  one  degree,  the  same  quantity  of 
heat  is  always  necessary. 

Now,  if  we  assume  that  the  constant  obtained  by 
multiplying  the  specific  heats  by  the  atomic  weights  is 
6.4,  which  is  about  the  average  of  the  different  values 
found,  then  it  is  plain  that,  if  we  divide  this  number  by 
the  specific  heat  of  an  element,  we  shall  obtain  a  number 
which  is  very  near  the  atomic  weight.  If  we  call  the 
atomic  weight  A,  and  the  specific  heat  H,  the  following 
equation  expresses  the  relation  : 


If  this  law  is  without  exceptions,  it  is  plain  that,  in  order 
to  determine  the  atomic  weight  of  an  element,  it  is  only 
necessary  to  determine  its  specific  heat,  and  divide  this 
into  6.4.  The  result  will  be  very  nearly  the  atomic 
weight.  Knowing  thus  very  nearly  what  the  atomic 
weight  is,  it  is  a  comparatively  simple  matter  to  deter- 
mine it  with  great  accuracy  by  means  of  chemical  analy- 
sis. Unfortunately  there  are  some  marked  exceptions 
to  the  law. 

Exceptions  to  the  Law  of  Specific  Heats.  —  The  elements 
glucinum,  carbon,  boron,  and  silicon  form  exceptions  to 
the  law  of  specific  heats  as  this  law  has  been  stated  above. 
At  ordinary  temperatures  they  do  not  follow  the  law.  As 
the  temperature  is  raised,  however,  the  specific  heat  of 
these  elements  changes  markedly,  until  finally,  in  the  cases 
of  carbon  and  silicon,  a  point  is  reached  beyond  which 
there  is  no  marked  change.  Thus,  at  600°  the  specific 
heat  of  diamond  is  0.441,  and  at  985°  it  is  0.449.  That 
of  silicon  is  0.201  at  185°,  and  0.203  at  332°.  At  these 


452  INORGANIC  CHEMISTRY. 

temperatures  the  elements  obey  the  law.  From  elabo- 
rate studies  which  have  been  made  on  this  subject,  it  ap- 
pears that  the  law  should  be  modified  to  read  as  follows  : 

The  specific  heats  of  the  elements  vary  with  the  tem- 
perature ;  but  for  every  element  there  is  a  temperature,  Tt 
above  which  variations  are  very  slight.  The  product  of 
the  atomic  weight  by  the  constant  value  of  the  specific 
heat  is  nearly  a  constant,  lying  between  5.5  and  6.5. 

Notwithstanding  the  irregularities  referred  to,  the  law 
of  specific  heats,  commonly  called,  from  the  discoverers, 
the  law  of  Dulong  and  Petit,  is  of  great  value  in  the  de- 
termination of  atomic  weights. 

Raoult's  Methods  for  the  Determination  of  Molecular 
Weights. — One  great  difficulty  encountered  in  the  study 
of  chemical  compounds  is  the  determination  of  the  mole- 
cular weights  of  those  which  are  not  gases  or  cannot  be 
converted  into  vapor  by  heat.  From  some  studies  on 
the  freezing-points  of  solutions,  it  appears  that  quantities 
of  compounds  proportional  to  their  molecular  weights 
cause  the  same  lowering  of  the  freezing-points,  provided 
the  solvent  does  not  act  chemically  upon  the  compound. 
This  fact  makes  it  possible  to  determine  the  molecular 
weights  of  substances  which  cannot  be  converted  into 
vapor,  but  which  can  be  dissolved.  The  application  of 
the  method  is  simple.  Suppose  water  to  be  the  solvent 
used.  We  know  that  this  liquid  solidifies  or  freezes  at 
0°.  Now,  it  is  found  that  by  dissolving  a  certain  quan- 
tity of  some  substance  in  a  certain  quantity  of  water  the 
freezing-point  is  lowered  say  .5°.  Further,  the  quantities 
of  other  substances  which  are  necessary  to  lower  the 
freezing-point  of  the  same  quantity  of  water  to  the  same 
extent  can  be  determined.  These  quantities  are  propor- 
tional to  the  molecular  weights  according  to  the  law  of 
Eaoult.  If,  therefore,  among  the  substances  studied 
there  is  one  the  molecular  weight  of  which  can  be  deter- 
mined by  the  method  of  Avogadro,  it  is  possible  to  de- 
termine the  molecular  weights  of  all  of  them  by  the 
method  of  Baoult,  as  will  readily  be  seen. 

So  also  it  has  been  shown  that  quantities  of  com- 
pounds proportional  to  their  molecular  weights  cause 


DISSOCIATION  OF  A  DISSOLVED  SUBSTANCE.      453 

the  same  raising  of  the  boiling-points,  provided  the 
solvent  does  not  act  chemically  upon  the  compounds  or 
cause  them  to  break  down  into  their  ions.  Convenient 
methods  have  been  devised  for  the  determination  of 
molecular  weights  of  dissolved  substances,  the  methods 
being  based  upon  observations  on  the  boiling-points 
and  freezing-points.  It  should  be  noted  that,  when  the 
molecular  weight  of  a  substance  in  solution  has  been 
determined,  it  does  not  follow  that  the  substance  has 
the  same  molecular  weight  when  in  the  solid  condi- 
tion. This  is  a  matter  in  regard  to  which  we  have  prac- 
tically no  knowledge.  It  is  quite  possible  that  the  mole- 
cules of  solid  substances  may  be  made  up  of  large 
aggregates  of  the  simple  molecules,  such  as  probably 
exist  in  solutions  or  in  vapors.  There  is,  however,  no 
method  at  present  known  that  makes  a  determination 
of  the  complexity  of  these  molecules,  or  molecular  aggre- 
gates, possible. 

Determination  of  the  Extent  of  Dissociation  of  a  Dis- 
solved Substance. — The  effect  upon  the  boiling-point  or 
freezing-point  of  a  solution  caused  by  the  presence  of  a 
dissolved  substance  is  proportional  to  the  number  of 
molecules  in  the  solution  or  the  number  of  individual 
particles,  whether  these  are  undecomposed  molecules  or 
the  ions  formed  as  a  result  of  dissociation.  Any  sub- 
stance that  is  dissociated  in  solution  will  give  abnormal 
results  if  the  attempt  is  made  to  determine  its  molecu- 
lar weight  by  observations  on  the  boiling-point  or  the 
freezing-point  of  its  solutions.  This  method  is  there- 
fore not  applicable  to  solutions  of  electrolytes.  On  the 
other  hand,  the  study  of  such  solutions  has  shown  that 
there  is  an  increased  lowering  of  the  freezing-point  for 
the  same  weight  of  solvent  as  the  dilution  becomes 
greater,  a  fact  that  points  clearly  to  the  conclusion  that 
as  the  solution  is  diluted  there  is  greater  and  greater 
dissociation,  and  advantage  can  be  taken  of  this  fact 
for  the  purpose  of  determining  the  extent  to  which  dis- 
sociation has  taken  place  in  a  solution  of  an  electrolyte 
in  water. 


454  INORGANIC  CHEMISTRY. 

In  general  only  organic  compounds  come  within  range 
of  the  methods  of  Raoult.  These  methods  are  now  ex- 
tensively used  in  the  study  of  the  compounds  of  carbon, 
simple  forms  of  apparatus  having  been  devised  for  this 
purpose.  The  recognition  of  the  fact  that  electrolytes 
do  not  obey  the  simple  law  that  holds  good  in  the  case 
of  non-electrolytes  led  Arrhenius  to  the  idea  that  the 
former  are  dissociated  in  solution — an  idea  which  has 
proved  of  great  service  to  the  science,  and  is  likely  to 
revolutionize  the  views  of  chemists  in  regard  to  the 
action  of  chemical  substances  upon  each  other  in  solu- 
tion. 


CHAPTER  XXIV. 

BASE-FORMING  ELEMENTS— GENERAL  CONSIDERATIONS 

Introductory. — The  elements  thus  far  considered  be- 
long for  the  most  part  to  the  class  of  acid-forming 
elements,  or  those  whose  compounds  with  oxygen  and 
hydrogen  have  acid  properties.  All  the  members  of 
Family  VII,  Group  B,  are  acid-forming,  while  the  single 
member  of  Group  A  of  the  same  family  is  both  acid- 
forming  and  base-forming.  All  the  members  of  Family 
VI,  Group  B,  are  acid-forming,  while  the  members  of 
Group  A  of  this  family  are  .both  acid-forming  and  base- 
forming.  In  Family  V,  Group  B,  there  is  observed  a 
gradation  of  properties,  the  group  beginning  with  strong- 
ly marked  acid-forming  elements  and  ending  with  an  ele- 
ment, bismuth,  which  is  more  basic  than  acid  in  char- 
acter. The  elements  of  Group  A,  Family  V,  are  both 
acid-forming  and  base-forming,  but  they  have  not  as 
sharply  marked  characteristics  as  the  elements  of  Fam- 
ilies VI  and  VII.  Passing  now  to  Family  IV,  we  found 
that  the  two  most  important  members,  carbon  and  sili- 
con, belong  to  Group  A.  These  two  elements  always  act 
as  acid-formers.  A  gradation  of  properties  is  observed 
in  passing  from  silicon  to  thorium.  The  members  of 
Group  B  of  this  family  have  the  properties  of  the  base- 
forming  elements  much  more  strongly  marked  than  those 
of  the  acid-formers.  There  are  still  four  families  to  be 
studied.  These  are  families  I,  II,  III,  and  VIII,  the 
members  of  which  are  almost  exclusively  base-forming 
elements.  The  compounds  of  these  elements  with  hy- 
drogen and  oxygen  are  bases,  or,  in  other  words,  have  the 
power  to  neutralize  acids.  Their  oxides  are  for  the  most 
part  basic.  An  exception  to  this  is  found  in  the  case  of 
boron,  already  considered,  which  forms  a  weak  acid — 
boric  acid.  Its  oxide  is  only  slightly  basic.  The  most 

(455) 


456  INORGANIC  CHEMISTRY. 

strongly  marked  examples  of  base-forming  elements  are 
those  which  occur  in  Family  I,  Group  A  ;  then  follow  in 
order  those  of  Group  A,  Family  II,  and  Group  A,  Fam- 
ily III.  The  resemblance  between  the  members  of 
Group  B,  Family  Iv  and  those  of  Group  A  of  the  same 
family  is  less  striking  than  the  resemblance  between  the 
two  groups  of  any  other  family.  Between  the  members 
of  Group  B,  Family  II,  and  those  of  Group  A  of  the  same 
family  there  is  a  general  resemblance,  while  there  are 
also  differences.  A  similar  remark  applies  to  the  rela- 
tions between  Groups  A  and  B,  Family  III.  The  mem- 
bers of  Family  VIII  occupy  a  somewhat  exceptional 
position,  as  has  already  been  pointed  out.  Each  group 
of  which  this  family  consists  is  made  up  of  three  very 
similar  elements  with  atomic  weights  which  differ  but 
little  from  one  another. 

Metallic  Properties. — It  has  long  been  customary  to 
divide  the  elements  into  two  classes — the  metals  and  the 
non-metals.  This  classification  was  originally  based  upon 
differences  in  the  physical  properties  of  the  elements,  the 
name  metal  being  applied  to  those  elements  which  have 
what  is  known  as  a  metallic  lustre,  are  opaque,  and  are 
good  conductors  of  heat  and  electricity.  All  those  ele- 
ments which  do  not  have  these  properties,  are  called  non- 
metals.  Gradually  the  name  metal  came  to  signify  an 
element  which  has  the  power  to  replace  the  hydrogen  of 
acids  and  form  salts,  and  the  name  non-metal  to  signify 
an  element  which  has  not  this  power.  This  classifica- 
tion, as  will  be  seen,  is  practically  the  same  as  that  which 
divides  the  elements  into  acid-forming  and  base-forming. 
The  latter  are  the  metals,  the  former  are  the  non-metals. 
The  imperfection  of  this  classification  has  already  been 
commented  upon,  the  imperfection  arising  from  the  fact 
that  some  elements  belong  to  both  classes. 

Order  in  which  the  Base-forming  Elements  will  be  Taken 
up. — In  studying  the  base-forming  elements,  it  appears 
best  to  begin  with  those  which  have  the  most  strongly 
marked  character.  These  are  the  members  of  Family  I, 
Group  A.  It  further  appears  best  to  adhere  as  closely 
as  possible  to  the  arrangement  in  the  periodic  system. 


OCCURRENCE  OF  THE  METALS.  457 

Accordingly,  the  following  order  will  be  observed  in  the 
presentation  of  the  elements  yet  to  be  studied : 

1.  Elements  of  Family  I,  Group  A,  or  the  Potassium 
Group,  consisting  of  lithium,  sodium,  potassium,  rubid- 
ium, and  caesium. 

2.  Elements  of  Family  II,  Group  A,  or  the   Calcium 
Group,    consisting    of    glucinum,   magnesium,    calcium, 
strontium,  barium,  and  erbium. 

3.  Elements  of  Family  III,  Group  A,  or  the  Aluminium 
Group,  consisting  of  aluminium,  scandium,  yttrium,  lan- 
thanum, and  ytterbium. 

4.  Elements   of   Family  I,  Group    B,  or   the    Copper 
Group,  consisting  of  copper,  silver,  and  gold. 

5.  Elements  of  Family  II,  Group  B,  or  the  Zinc  Group, 
consisting  of  zinc,  cadmium,  and  mercury. 

6.  Elements  of  Family  III,  Group  B,  or  the  Gallium 
Group,  consisting  of  gallium,  indium,  and  thallium. 

7.  Elements  of  Family  IY,  Group  B,  or  the  Tin  Group, 
consisting  of  germanium,  tin,  and  lead. 

8.  Elements  of  Family  Y,  Group  A,  or  the  Vanadium 
Group,  consisting  of  vanadium,  columbium,  didymium, 
and  tantalum. 

9.  Elements  of  Family  VI,  Group  A,  or  the  Chromium 
Group,  consisting  of  chromium,  molybdenum,  tungsten, 
and  uranium. 

10.  Elements  of  Family  VII,  Group  A,  or  the  Manganese 
Group,  of  which  manganese  is  the  only  representative. 

11.  Elements    of    Family  VIII,   of  which   there   are 
three  groups  : 

(A)  The  Iron   Group,  consisting  of  iron,  nickel,  and 
<?obalt ; 

(B)  The  Palladium    Group,    consisting  of   ruthenium, 
rhodium,  and  palladium  ;  and 

(C)  The  Platinum  Group,  consisting  of  osmium,  irid- 
ium,  and  platinum. 

Occurrence  of  the  Metals. — One  of  the  first  questions 
that  suggests  itself  in  connection  with  each  element  is, 
In  what  forms  of  combination  does  it  occur  in  nature  ? 
The  chemical  compounds  which  occur  ready-formed  in 
nature  are  called  minerals;  and  the  minerals,  and  mixtures 


458  INORGANIC  CHEMISTRY. 

of  minerals,  from  which  the  metals  are  extracted  for  prac- 
tical purposes  are  called  ores.  The  most  common  ores 
are  oxides  and  sulphides.  Examples  of  these  are  the 
ores  of  iron,  tin,  copper,  lead,  and  zinc.  The  carbonates 
also  occur  in  large  quantity  in  nature,  and  are  used  for 
the  purpose  of  preparing  some  of  the  metals.  The  car- 
bonate of  zinc,  for  example,  is  a  valuable  ore  of  zinc. 

Extraction  of  the  Metals  from  their  Ores. — The  detailed 
study  of  the  methods  used  in  the  extraction  of  the  metals 
from  their  ores  is  the  object  of  metallurgy.  Besides  the 
methods  used  on  the  large  scale,  there  are  others  which  are 
only  used  in  the  laboratory.  The  most  common  method 
of  extracting  metals  from  their  ores  is  that  used  in  'the 
case  of  iron,  which  consists  in  heating  the  oxides  with  char- 
coal. If  the  ores  used  are  not  oxides,  they  must  first  be 
converted  into  oxides  before  this  method  is  applicable. 
This  can  generally  be  accomplished  by  heating  the  ores 
in  contact  with  the  air.  Under  these  circumstances  the 
natural  carbonates,  sulphides,  and  hydroxides,  are  con- 
verted into  oxides.  These  changes  are  illustrated  by  the 
following  equations  : 

FeC03  =  FeO  +  CO2 ; 
2FeO  +  O  =  Fe2O3 ; 
2FeS2  -f  HO  =  Fe2O3  +  4SO2 ; 
2Fe(OH)3  =  Fe2O3  +  3H2O. 

A  second  method  consists  in  reducing  the  oxide  by  heat- 
ing it  in  a  current  of  hydrogen.  This  has  been  illustrated 
in  the  action  of  hydrogen  upon  copper  oxide,  when  the 
following  reaction  takes  place  : 

CuO  +  H2  =  H20  +  Cu. 

The  method  is  efficient  for  many  oxides,  but  is  expen- 
sive and  is  not  used  on  the  large  scale. 

Another  method  of  extraction  consists  in  treating  the 
chloride  of  a  metal  with  sodium.  This  is  illustrated  in 
the  preparation  of  magnesium,  which  is  made  by  heating 
together  magnesium  chloride  and  sodium : 

MgCl2  +  2Na  =  2NaCl  +  Mg. 


COMPOUNDS  OF  THE  METALS.  459 

Such  a  method  is  employed  only  in  case  it  is  impossible 
or  extremely  difficult  to  reduce  the  oxide. 

Besides  the  above  methods,  there  are  others  which  will 
be  described  under  the  individual  metals. 

The  Properties  of  the  Metals. — As  we  shall  find,  the 
metals  differ  very  markedly  from  one  another.  Some 
are  light,  floating  on  water,  as  lithium,  sodium,  etc. ; 
some  are  extremely  heavy,  as  lead,  platinum,  etc.  Some 
combine  with  oxygen  with  great  energy  ;  others  form  very 
unstable  compounds  with  oxygen.  Some  form  strong 
bases ;  others  form  weak  bases.  In  general,  those  ele- 
ments which  are  lightest,  or  which  have  the  lowest 
specific  gravity,  are  the  most  active  chemically,  while 
those  which  have  the  highest  specific  gravity  are  the 
least  active.  Among  the  former  are  lithium,  sodium,  and 
potassium  ;  among  the  latter  are  lead,  gold,  and  platinum. 

Compounds  of  the  Metals. — The  principal  compounds 
of  the  metals  may  be  conveniently  classified  as : 

a.  Compounds  with  fluorine,  chlorine,   bromine,  and 
iodine  ;  or  the  fluorides,  chlorides,  bromides,  and  iodides. 

b.  Compounds  with  oxygen,  and  with  oxygen  and  hy- 
drogen ;  or  the  oxides  and  hydroxides. 

c.  Compounds  with  sulphur,  and  with  sulphur  and  hy- 
drogen ;  or  the  sulphides  and  hydrosulphides. 

d.  Compounds  with  nitrogen  ;  or  the  nitrides. 

e.  Compounds  with  carbon  and  with  silicon ;  or  the 
carbides  and  silicides. 

f.  Compounds  with  the  acids  of  nitrogen  ;  or  the  ni- 
trates and  nitrites. 

g.  Compounds  with  the  acids  of  chlorine,  bromine,  and 
iodine  ;  or  the  chlorates,  bromates,  iodates,  hypochlorites,  etc. 

h.  Compounds  with  the  acids  of  sulphur,  selenium, 
and  tellurium  ;  or  the  sulphates,  sulphites,  etc. 

i.  Compounds  with  carbonic  acid  ;  or  the  carbonates. 

j.  Compounds  with  the  acids  of  phosphorus,  arsenic, 
and  antimony  ;  or  the  phosphates,  arsenates,  etc. 

k.  Compounds  with  silicic  acid  ;  or  the  silicates. 

1.  Compounds  with  boric  acid  ;  or  the  borates. 

It  is  more  important  to  become  acquainted  with  the 
general  methods  of  preparation  and  the  general  properties 


460  INORGANIC  CHEMISTRY. 

of  the  more  important  compounds  than  to  learn  details 
concerning  many  individual  members  of  each  class.  Only 
those  compounds  which  illustrate  general  principles, 
or  which,  owing  to  some  application,  happen  to  be  of 
special  interest,  need  be  fully  treated  in  this  book. 

The  acids  of  which  the  salts  are  derivatives  are  already 
known  to  us,  and  in  dealing  with  the  acids  frequent 
reference  has  been  made  to  the  methods  of  making  the 
salts,  and  to  some  of  their  more  important  properties. 
It  will  be  well,  before  taking  up  the  metals  systemati- 
cally, briefly  to  treat  of  the  general  methods  of  prepara- 
tion, and  the  general  properties  of  the  different  classes 
of  metallic  compounds.  It  must  be  borne  in  mind,  how- 
ever, that  the  only  way  to  become  familiar  with  these 
substances  and  their  relations  is  by  working  with  them  in 
the  laboratory. 

Chlorides. — The  chlorides,  as  well  as  the  fluorides, 
bromides,  and  iodides,  may  be  regarded  as  the  salts  of 
hydrochloric,  hydrofluoric,  hydrobromic,  and  hydriodic 
acids,  or  simply  as  compounds  of  the  metals  with  the 
members  of  the  chlorine  family.  The  most  important 
of  these  compounds  are  the  chlorides,  and  these  well 
illustrate  the  conduct  of  the  others. 

The  chlorides  are  made  by  treating  a  metal  with  chlo- 
rine, or  with  hydrochloric  acid ;  by  treating  an  oxide 
or  a  hydroxide  with  hydrochloric  acid ;  by  treating  an 
oxide  with  chlorine  and  a  reducing  agent,  like  carbon ; 
by  treating  a  salt  of  a  volatile  acid  with  hydrochloric 
acid  ;  by  treating  a  salt  of  an  insoluble  acid  with  hydro- 
chloric acid ;  by  adding  hydrochloric  acid  or  a  soluble 
chloride  to  a  solution  containing  a  metal  with  which 
chlorine  forms  an  insoluble  compound ;  and  by  adding  to 
a  solution  of  a  chloride  a  salt,  the  acid  of  which  forms 
with  the  metal  of  the  chloride  an  insoluble  salt,  while 
the  metal  contained  in  it  forms  with  chlorine  a  soluble 
chloride. 

Only  two  of  the  above  methods  are  peculiar  to  chlo- 
rides. These  are  the  treatment  of  the  metals  with  chlo- 
rine, and  the  treatment  of  oxides  with  chlorine  and  a 
reducing  agent.  The  others  involve  principles  which 


CHLORIDES.  461 

are  also  involved  in  the  preparation  of  all  salts,  and  they 
may  therefore  be  treated  of  in  a  general  way. 

The  formation  of  chlorides  by  direct  treatment  of  the 
metals  with  chlorine  is  the  simplest  method  of  all.  It 
has  been  illustrated  in  studying  chlorine.  It  was  found 
that  chlorine  combines  with  other  elements  with  great 
ease.  Thus,  iron,  copper,  and  tin  combine  with  it,  as 
represented  in  the  following  equations  : 

Cu  +01,    =  CuCl2; 
2Fe  +  3C12  =  2FeCl3  ; 
Sn 


The  preparation  of  chlorides  by  treating  oxides  with 
chlorine  and  a  reducing  agent  has  been  illustrated  in 
the  making  of  boron  trichloride  and  of  silicon  tetra- 
chloride.  It  is  used  in  making  aluminium  trichloride. 
For  this  purpose,  chlorine  is  passed  over  a  heated  mix- 
ture of  aluminium  oxide  and  charcoal,  when  reaction 
takes  place  according  to  the  following  equation  : 

A1203  +  3C  +  3C12  =  2A1C13  +  3CO. 

The  interesting  character  of  this  reaction  was  referred  to 
in  connection  with  the  similar  preparation  of  the  chlo- 
rides of  boron  and  silicon.  In  this  case,  as  in  those, 
there  are  two  reactions  involved.  The  carbon  alone  can- 
not reduce  the  oxide  ;  nor  can  the  chlorine  alone  decom- 
pose it  to  form  the  chloride.  But  when  the  carbon  and 
chlorine  act  together,  they  assist  each  other,  and  as  a 
consequence  the  oxide  is  transformed  into  the  chloride. 

The  other  methods  for  preparing  chlorides  are,  as  has 
been  said,  general  in  character  and  are  applicable  to 
most  salts. 

Formation  of  Salts  in  General. 

1.  By  treating  a  metal  with  an  acid.  —  This  is  the  sim- 
plest method.  It  has  been  illustrated  in  the  preparation 
of  zinc  sulphate  by  the  action  of  zinc  on  sulphuric  acid  : 

Zn  +  H2SO4  =  ZnS04  +  H2. 


462  INORGANIC  CHEMISTRY. 

Other  common  examples  are  those  represented  in  the 
follow  equations : 

Fe  +  H2S04  =  FeS04  +  Ha ; 
Zn+2HCl    =  ZnCla  +  H3. 

2.  By  treating  an  oxide  or  a  hydroxide  with  an  acid. — 
This  is  of  more  general  application  than  the  preceding 
method.     As  it  has  been  studied  in  some  detail  in  con- 
nection with  the  subject  of  salts  (see  pp.  129-133),  it  need 
not  be  further  considered  here. 

3.  By  treating  the  salt  of  a  volatile  acid  with  another  acid. 
— This  method  has  been  repeatedly  illustrated  in  the  de- 
composition of  carbonates  and  nitrites  by  acids  in  gen- 
eral.    While  carbonic  acid  and  nitrous  acid  themselves- 
are  perhaps  not  formed  in  these  reactions,  and  we  can- 
not say  that  the  carbonates  and  nitrites  are  salts  of  vol- 
atile acids,  yet  the  decomposition-products  of  these  acids 
are  volatile  at  ordinary  temperatures.      The  decompo- 
sition of   carbonates   by   acids    has   been   pretty  fully 
studied,  though  attention  was  not  directed  to  the  fact 
that  this  kind  of  action  may  be  utilized  for  the  purpose 
of  making  salts.      As  some  carbonates  occur  in  large 
quantity  in  nature  or  in  the  market,  salts  are  frequently 
made  by  treating  them  with  acids.     Thus,  magnesium 
sulphate  is  made  by  treating  magnesium  carbonate  with 
sulphuric  acid : 

MgC03  +  H,SO,  =  MgSO,  +  H,0  +  CO, ; 

and  calcium  chloride  is  made  by  dissolving  calcium  car- 
bonate in  hydrochloric  acid : 

CaCO,  +  2HC1  =  CaCl2  +  H2O  +  CO2 ;  etc. 

4.  By  treating  a  salt  of  an  insoluble  acid  with  another 
acid. — This  case  does  not  occur  practically,  as  there  are 
no  common,  insoluble  acids.     The  principle  involved  is 
illustrated  to  some  extent  by  the  decomposition  of  a  sol- 
uble silicate.     Sodium  silicate,  for  example,  is  soluble. 
When   its   solution  in   water   is   treated   with   an   acid 


GENERAL  PROPERTIES  OF  THE  CHLORIDES.       463 
the  silicic  acid  is  partly  precipitated,  as  we  have  seen 

Na2SiO3  +  2HC1  +  HaO  =  Si(OH)4  +  2NaCl. 

The  silicic  acid  formed  is,  however,  not  perfectly  in- 
soluble in  water,  so  that  the  reaction  is  not  complete. 
In  any  case  the  reaction  is  not  one  that  is  used  for  the 
preparation  of  salts. 

5.  By   the  action  of  two  salts  upon  each  other. — This 
method  can  be  best  described  by  means  of  an  example. 
Suppose  it  is  desired  to  prepare  copper  chloride  by  the 
action  of  two  salts  upon  each  other.     Copper  chloride 
is  soluble.     If  copper  sulphate  and  barium  chloride  are 
brought  together  in  solution,  the  products  are  insoluble 
barium  sulphate  and  soluble  copper  chloride  : 

CuSO4  +  BaCl2  =  BaS04  +  CuCl2. 

By  simply  filtering  off  the  barium  sulphate,  a  solution 
of  copper  chloride  is  obtained. 

6.  By  precipitation. — This  method  is  illustrated  in  the 
formation  of  barium  sulphate,  referred  to  in  the  last  par- 
agraph.     Obviously,  it  is  applicable  only  to  difficultly 
soluble  or  insoluble  salts.     Many  carbonates  and  phos- 
phates can  be  made  in  this  way. 

General  Properties  of  the  Chlorides. — Most  of  the  chlo- 
rides of  the  metals  are  soluble  in  water  without  decom- 
position, though  many  of  them  are  decomposed  when 
heated  to  a  sufficiently  high  temperature  with  water.  It 
will  be  remembered  that  the  chlorides  of  the  non-metal- 
lic or  acid-forming  elements  are  decomposed  by  water, 
yielding  the  corresponding  oxides  or  hydroxides.  The 
chlorides  of  some  elements  which  are  partly  basic  and 
partly  acid  are  only  partly  decomposed.  This  is  illus- 
trated by  the  chloride  of  antimony,  which  with  water 
forms  an  oxy chloride  : 

SbCl3  +  H20  =  SbOCl  +  2HC1. 


464  INORGANIC  CHEMISTRY. 

The  chlorides  of  the  most  strongly  marked  metals, 
like  potassium,  sodium,  etc.,  are  not  decomposed  by 
water.  Calcium  chloride  dissolves  with  great  ease,  and, 
if  the  solution  is  evaporated,  the  chloride  is  again 
obtained.  If,  however,  the  attempt  is  made  to  drive  off 
all  the  water  by  heat,  some  of  the  chloride  is  converted 
into  the  oxide  as  represented  in  the  equation 

CaCl,  +  H20  =  CaO  +  2HC1. 

Magnesium  chloride  is  completely  decomposed,  if  its 
solution  in  water  is  evaporated  to  dryness,  the  action 
being  the  same  in  character  as  that  which  takes  place  in 
the  case  of  calcium  chloride.  The  chlorides  of  iron  and 
aluminium  and  of  many  other  metals  act  in  the  same  way. 
Silver  chloride  and  mercurous  chloride,  HgCl,  are  insol- 
uble in  water.  Lead  chloride  is  difficultly  soluble  in 
water.  If,  therefore,  on  adding  hydrochloric  acid  or  a 
soluble  chloride  to  a  solution,  a  precipitate  is  formed, 
the  conclusion  is  generally  justified  that  one  or  more  "of 
the  three  metals — silver,  lead,  or  mercury — is  present. 
By  taking  into  account  the  differences  between  these 
chlorides,  it  is  not  difficult  to  decide  of  which  of  them  a 
precipitate  consists. 

The  chlorides  are  for  the  most  part  stable  when  heated, 
though  a  few  lose  some  of  their  chlorine  just  as  phos- 
phorus pentachloride  does.  An  example  of  this  is  pre- 
sented by  platinic  chloride,  PtCl4,  which  when  heated 
breaks  down  into  platinous  chloride,  PtCl2,  and  chlorine  : 

PtCl4  =  PtCl,  +  C12. 

The  chlorides  are  for  the  most  part  decomposed  when 
treated  with  sulphuric  acid,  as  has  been  shown  in  the 
action  of  sulphuric  acid  upon  sodium  chloride.  Under 
these  circumstances  hydrochloric  acid  is  given  off,  and 
the  sulphate  of  the  metal  with  which  the  chlorine  was 
in  combination  is  formed.  In  general,  the  reaction  is 
represented  by  such  equations  as  the  following : 

2MC1  +  H2S04  =  M2S04  +  2HC1 ; 
MC12  +  H2S04  =  MS04  +  2HC1 ;  etc. 


THE  SO-CALLED  DOUBLE  CHLORIDES.  465 

Under  ordinary  circumstances,  chlorides  are  not  decom- 
posed by  any  acid  except  sulphuric  acid. 

The  So-called  Double  Chlorides  and  Similar  Compounds 
of  Fluorine,  Bromine,  and  Iodine. — These  compounds  and 
their  relations  to  the  oxygen  salts  have  been  repeatedly 
referred  to.  Many  chlorides  combine  with  the  chlorides 
of  the  stronger  metals,  like  sodium  and  potassium,  forming 
well-characterized  compounds.  Generally,  these  double 
chlorides  are  analogous  to  the  oxygen  salts  in  com- 
position, differing  from  them  only  by  containing  two 
atoms  of  chlorine  in  the  place  of  each  of  the  oxygen 
atoms.  As  examples  of  these  salts  of  the  Moro-acids 
those  which  are  formed  by  the  chlorides  of  platinum, 
antimony,  chromium,  and  gold  may  be  mentioned. 
Platinic  chloride,  PtCl4,  combines  with  other  chlorides, 
forming  salts  of  the  general  composition  expressed  by  the 
formula  PtCl4  +  2MC1,  or  M2PtCl6.  Antimony  chlo- 
ride, combines  with  three  molecules  of  potassium  chlo- 
ride forming  the  compound  Sb013  +  3KC1,  or  K3SbCl6. 
Chromium  chloride  forms  similar  compounds,  K3CrCl6, 
Na3CrCl6 ;  and  gold  chloride  forms  compounds  of  the 
general  formula  MAuCl4,  which  may  be  regarded  as 
made  up  of  one  molecule  of  auric  chloride,  AuCl3,  and 
one  molecule  of  a  chloride  like  potassium  chloride. 
A  careful  study  of  the  double  chlorides  and  the '  similar 
compounds  of  fluorine,  bromine,  and  iodine  shows 
that  the  chlorides  of  sodium  and  potassium,  and  of  the 
other  elements  of  the  group  to  which  these  metals  be- 
long, combine  with  most  other  chlorides  to  form  so- 
called  double  salts,  and  that  the  number  of  molecules  of 
potassium  or  sodium  chloride  which  combine  with  another 
chloride  is  limited  by  the  number  of  chlorine  atoms  contained 
in  the  other  chloride.*  Thus,  a  chloride  of  the  formula 
MC12  may  form  the  double  chlorides  MC12.KC1  and 
MC12.2KC1,  but  not  MC12.3KC1.  So,  further,  a  chloride  of 
the  formula  MC13  may  form  three  different  double  chlo- 
rides with  the  same  metallic  chloride.  Those  with  potas- 
sium will  have  the  formulas  MC13.KC1,  MC13.2KC1,  and 

*  There  are  a  few  exceptions  to  this  rule,  but  it  undoubtedly  holds 
good  in  by  far  the  largest  number  of  cases. 


46G  INORGANIC  CHEMISTRY. 

MC13.3KC1,  but  a  double  chloride  of  the  formula 
MC13.4KC1  and  more  complicated  cases  seem  to  be  im- 
possible. Double  fluorides  are  known  in  large  numbers. 
Among  the  best-known  are  the  fluosilicates.  Aluminium 
forms  double  fluorides,  one  of  which,  having  the  formula 
Na3AlF6  or  AlF3.3NaF,  is  the  well-known  mineral 
cryolite.  All  these  so-called  "  double  salts "  are  easily 
explained  by  the  aid  of  the  hypothesis  that  the  halogen 
contained  in  them  has  a  valence  greater  than  one,  and 
that  a  double  atom,  like  C12,  F2,  etc.,  or  -C1-C1-,  -F-F-, 
plays  the  same  part  that  oxygen  does  in  the  oxygen 
salts.  The  following  table  contains  the  general  for- 
mulas of  the  possible  double  chlorides  with  potassium 
chloride,  according  to  the  above  view  concerning  them : 

MC1.KC1    MC12.KC1      MC13.KC1      MC14.KC1 
MC12.2KC1    MC13.2KC1    MC14.2KC1 
MC13.3KC1    MC14.3KC1 
MC14.4KC1 

These  may  also  be  written  thus  : 

KMC1,    KMC13     KMC14     KMC15 

K2MC14  K2MC16  K2MC16 

K3MC16  K3MC17 

K4MC18 

Different  Chlorides  of  the  Same  Metal. — Just  as  sul- 
phur, selenium,  phosphorus,  and  the  other  acid-forming 
elements  combine  with  chlorine  and  the  other  members 
of  the  chlorine  group  in  more  than  one  proportion,  so 
many  of  the  metals  combine  with  the  members  of  the 
chlorine  group  in  more  than  one  proportion.  Thus,  mer- 
cury forms  the  two  chlorides,  HgCl2  and  HgCl,  known 
respectively  as  mercuric  and  mercurous  chlorides ;  iron 
forms  the  two  chlorides  FeCl3  and  FeCl2,  known  as  ferric 
and  ferrous  chlorides  ;  and  tin  forms  stannic  chloride, 
SnCl4,  and  stannous  chloride,  SnCl2.  The  conversion  of 
a  higher  chloride  into  a  lower  one  is  called  an  act  of 


OXIDES.  467 

reduction.  The  change  can  generally  be  effected  by 
means  of  nascent  hydrogen : 

SnCl4  +  2H  =  SnCl2  +  2HC1. 
FeCl3  +    H  =  FeCla  +    HC1. 

The  conversion  of  a  lower  chloride  into  a  higher  one  is 
generally  spoken  of  as  an  act  of  oxidation,  for  the  reason 
that  it  is  most  commonly  effected  by  the  action  of  oxygen. 
Thus  the  most  convenient  way  to  transform  ferrous  chlo- 
ride into  ferric  chloride  is  to  treat  it  in  solution  in  hydro- 
chloric acid  with  an  oxidizing  agent,  when  a  double  action 
takes  place,  as  represented  in  the  following  equation  : 

2FeCl2  +  2HC1  +  O  =  2FeCl3  +  H2O. 

The  same  change  can  be  effected  by  the  direct  action  of 
chlorine : 

Fed,  +  01  =  FeCl3. 

In  this  case  it  would  obviously  be  incorrect  to  speak  of 
the  process  as  one  of  oxidation. 

Another  method  of  reduction,  besides  that  referred  to 
above,  involving  the  action  of  nascent  hydrogen,  is  that 
illustrated  in  the  equation 

2HgCl2  +  SnCl2  =  2HgCl  +  SnCl4. 

In  this  case  mercuric  chloride  is  changed  to  mercurous 
chloride  by  the  action  of  stannous  chloride.  The  latter 
unites  with  chlorine  so  readily  that  it  extracts  it  from 
some  other  chlorides,  and  is  itself  transformed  into 
stannic  chloride.  While,  therefore,  we  say  that  the 
stannous  chloride  reduces  the  mercuric  chloride,  it  is 
equally  true  to  say  that  the  mercuric  chloride  chlorinates 
the  stannous  chloride. 

Oxides. — The  oxides  occur  very  extensively  in  nature, 
and  are  among  the  most  common  ores  of  some  of  the 
important  metals.  The  oxides  of  iron,  tin,  and  man- 
ganese, for  example,  occur  in  nature.  They  can  be  made 
by  oxidizing  the  metals,  by  heating  nitrates,  carbonates, 
and  hydroxides,  and  by  heating  some  sulphides  in  con- 
tact with  the  air. 


468  INORGANIC  CHEMISTRY. 

When  magnesium  is  burned  it  is  converted  into  mag- 
nesium oxide : 

Mg  +  O  =  MgO. 

When  lead  nitrate  is  heated  it  gives  off  oxygen  and  an 
oxide  of  nitrogen,  and  lead  oxide  is  left  behind : 

Pb(NOs)3  =  PbO  +  2N02  +  O. 

When  calcium  carbonate  is  heated  it  yields  calcium  oxide 
and  carbon  dioxide : 

CaCO3  =  CaO  +  C02. 

When  aluminium  hydroxide,  Al(Ofi),,  is  heated  it  loses 
water,  and  aluminium  oxide  is  left  behind : 

2A1(OH)3  =  A12O3  +  3H2O. 

The  sulphide  of  iron,  when  heated  in  contact  with  the 
air,  or  "  roasted,"  is  converted  into  ferric  oxide  and  sul- 
phur dioxide. 

Most  of  the  oxides  of  the  metals  are  insoluble  in  water. 
Those  of  Group  A,  Family  I,  are  soluble,  but  are  con- 
verted by  water  into  the  corresponding  hydroxides. 

The  oxides  are  acted  upon  generally  by  acids  forming 
the  corresponding  salts.  If  the  salt  with  a  certain  acid 
is  insoluble,  the  salt  is  not  formed  by  the  action  of  that 
acid  on  the  oxide  unless  the  acid  or  its  anhydride  is 
fusible  and  not  volatile,  when  by  fusing  them  together  the 
salt  is  formed. 

Different  Oxides  of  the  Same  Metal. — Just  as  there  are 
different  chlorides  of  the  same  metal,  so  there  are  differ- 
ent oxides,  and  indeed  there  is  greater  variety  among 
these  than  among  the  chlorides.  Iron  forms  three  oxides, 
ferric  oxide,  Fe2O3,  ferroso-ferric  oxide,  Fe3O4,  and  ferrous 
oxide,  FeO ;  mercury  forms  the  two  oxides  HgO  and  Hg2O  ; 
etc.  The  lower  oxides  are  converted  into  the  higher 
by  oxidation,  and  the  higher  into  the  lower  by  reduc- 
tion. The  higher  oxides  of  several  of  the  metals  are 
acidic.  This  is  markedly  so  in  the  case  of  chromium  and 
manganese. 


HYDROXIDES.  469 

Hydroxides. — The  hydroxides  are  formed  by  treating 
oxides  with  water  and  by  decomposing  salts  by  adding 
soluble  hydroxides  to  their  solutions.  In  general,  when- 
ever a  salt  is  decomposed  by  a  strong  base,  the  base  of 
the  salt  separates  in  the  form  of  the  hydroxide.  The 
formation  of  a  hydroxide  by  the  action  of  water  on  an 
oxide  is  well  illustrated  by  the  action  of  water  on  lime 
or  calcium  oxide,  a  process  which  is  familiarly  known  as 
slaking : 

CaO  +  H2O  =t  Ca(OH),. 

Most  of  the  hydroxides  of  the  metals  are  insoluble  in 
water.  If  a  soluble  hydroxide  is  added  to  a  solution 
containing  a  metal  whose  hydroxide  is  insoluble,  the 
latter  is  precipitated.  Thus,  if  a  solution  of  sodium  hy- 
droxide is  added  to  a  solution  of  a  magnesium  salt, 
magnesium  hydroxide  is  precipitated  : 

MgSO4  +  2NaOH  =  Na2SO4  +  Mg(OH)2. 

So,  also,  when  a  solution  of  a  ferric  salt  is  treated  with 
sodium  hydroxide,  a  precipitate  of  ferric  hydroxide  is 
formed : 

FeCl3  +  3NaOH  =  3NaCl  +  Fe(OH)3. 

Only  the  hydroxides  of  the  members  of  the  potassium 
family,  and  some  of  the  members  of  the  calcium  family, 
are  soluble  in  water.  The  hydroxides  of  sodium  and 
potassium  are  called  alkalies.  The  solution  of  ammonia 
in  water  acts  like  a  soluble  hydroxide,  and  probably  con- 
tains ammonium  hydroxide,  NH4(OH),  formed  by  the 
action  of  water  on  ammonia  : 

NH3  +  H20  =  NH4(OH). 

Now,  when  any  one  of  the  soluble  hydroxides  is  added  to 
a  salt  containing  any  metal  which  does  not  belong  to  the 
potassium  or  calcium  family,  an  insoluble  compound  is 
formed. 

Decomposition  of  Salts  by  Bases. — The  decomposition 
of  salts  by  bases  is  analogous  to  the  decomposition  by 


470  INORGANIC  CHEMISTRY. 

acids.     When  a  soluble  base  acts  upon  a  salt,  there  are 
four  possible  kinds  of  action : 

1.  The  base  from  which  the  salt  is  derived  may  be 
volatile,  or  may  break  up,  yielding  a  volatile  product. 

In  this  case,  decomposition  takes  place  and  the  volatile 
base  is  given  off.  This  is  not  a  common  case  except 
among  the  compounds  of  carbon.  The  one  illustration 
which  we  have  had  is  the  decomposition  of  ammonium 
salts  by  calcium  hydroxide  and  sodium  hydroxide,  when 
the  volatile  compound  ammonia,  NH3,  is  given  off. 

2.  The  hydroxide,  or  base  from  which  the  salt  is  de- 
rived, may  be  insoluble  or  difficultly  soluble  in  water, 
and  not  volatile. 

In  this  case,  if  both  the  salt  and  the  base  are  in  solu- 
tion, decomposition  takes  place,  and  the  insoluble  or 
difficultly  soluble  hydroxide,  or  base,  is  precipitated. 
This  has  already  been  illustrated. 

3.  The  base  from  which  the  salt  is  derived  may  be 
soluble  and  not  volatile. 

This  is  the  case,  for  example,  when  sodium  hydroxide 
is  added  to  a  solution  of  potassium  nitrate.  Here  sodium 
nitrate,  potassium  nitrate,  sodium  hydroxide,  and  potas- 
sium hydroxide  may  all  be  present  in  the  solution,  and 
investigation  has  shown  that  all  are  present  and  that  the 
quantity  of  each  depends  upon  the  masses  of  the  sub- 
stances brought  together,  and  upon  their  affinities. 

4.  The  fourth  case  is  that  in  which  a  soluble  hydroxide 
forms  an  insoluble  salt  with  the  acid  of  a  soluble  salt, 
leaving  a  soluble  hydroxide  in  solution. 

This  is  illustrated  by  the  action  of  calcium  hydroxide 
on  a  solution  of  sodium  carbonate,  when  insoluble  cal- 
cium carbonate  is  thrown  down,  and  sodium  hydroxide 
remains  in  solution,  as  represented  in  the  equation 

Na2CO3  +  Ca(OH)2  =  2NaOH  +  CaCO3. 

Some  basic  hydroxides,  which  are  precipitated  by  solu- 
ble hydroxides,  have  a  weak  acid  character,  and,  after  they 
are  precipitated,  they  redissolve  in  an  excess  of  the  solu- 
ble hydroxide.  This  is  true,  for  example,  of  aluminium, 


DECOMPOSITION  OF  SALTS  BY  BASES.  471 

chromium,  and  lead.  The  salt-like  compounds  thus 
formed  are  generally  quite  unstable.  The  precipitation 
and  subsequent  solution  of  the  hydroxides  of  the  three 
metals  named  take  place  thus  : 

A1C1,        +  3NaOH  =  Al(OH),     +  3NaCl; 
A1(OH)3   +  3NaOH  =  Al(ONa)3  +  3HaO  ; 
CrCl3        +  3NaOH  =  Cr(OH),    +  3NaCl  ; 
Cr(OH)3  +  3NaOH  =  Cr(ONa)3  +  3H0O  ; 
Pb(NO3)2  +  2NaOH  =  Pb(OH)2  +  2NaNO3  ; 
Pb(OH)2  +  2NaOH  =  Pb(ONa)2  +  2H2O. 

In  some  cases  where  a  soluble  hydroxide  is  added  to  a 
salt,  an  oxide  is  precipitated  instead  of  the  hydroxide. 
This  is  analogous  to  the  formation  of  an  anhydride  of 
an  acid  instead  of  the  acid  itself,  as  when  carbonates, 
sulphites,  and  nitrites  are  decomposed. 

When  a  silver  salt  is  treated  with  a  soluble  hydroxide, 
silver  oxide  is  at  once  precipitated.  The  same  is  true  of 
mercury  salts  : 

2AgN03  +  2KOH  =  Ag20  +  H2O  +  2KNO3  ; 
HgCl2  +  2NaOH  =  HgO  +  H2O  +  2NaCl. 

It  is  probable  that  the  first  product  is  the  hydroxide, 
and  that  this  breaks  down  into  the  oxide  and  water  : 

2AgNO3  +  2KOH  =  2AgOH  +  2KNO3  ; 

2AgOH  =  Ag20  +  H20; 

HgCl2  +  2NaOH  =  Hg(OH)2  +  2NaCl  ; 


Some  hydroxides  are  converted  into  the  oxides  by 
simply  boiling  the  liquids  in  which  they  are  suspended. 
Thus,  when  a  salt  of  copper  is  treated  with  a  soluble  hy~ 
droxide,  copper  hydroxide  is  first  precipitated  ;  but  if  the 
solution  in  which  it  is  suspended  is  boiled,  it  is  soon 
changed  to  the  oxide  : 

CuSO4  +  2NaOH  =  Na2SO4  +  Cu(OH)2  ; 


472  INORGANIC  CHEMISTRY. 

The  hydroxides  corresponding  to  some  of  the  highei 
oxides  of  the  metals,  as  those  of  chromium  and  mangan- 
ese, are  acids. 

The  hydroxides  of  most  of  the  metals  are  decomposed 
by  heat  into  water  and  the  corresponding  oxides.  Those 
of  the  alkali  metals,  as  potassium  and  sodium,  are  not, 
however,  decomposed  by  heat. 

Sulphides.  —  Many  sulphides  are  found  in  nature,  as, 
for  example,  iron  pyrites,  FeS2  ;  lead  sulphide,  or  galen- 
ite,  PbS  ;  copper  pyrites,  FeCuS2  ;  etc.  They  are  made 
in  the  laboratory  by  heating  metals  with  sulphur  ;  by 
treating  solutions  of  salts  with  hydrogen  sulphide  ;  by 
treating  solutions  of  salts  with  soluble  sulphides  ;  and 
by  reducing  sulphates.  Attention  has  been  called  to  the 
fact  that  the  sulphides  are  analogous  in  composition  to 
the  oxides,  and  that  they  are  to  be  regarded  as  salts  of 
hydrogen  sulphide  formed  by  replacing  the  hydrogen  of 
the  acid  by  metals. 

The  formation  of  sulphides  by  the  direct  combination 
of  sulphur  with  the  metals  is  shown  in  the  formation  of 
lead  sulphide  and  copper  sulphide  : 


2Cu  +  S  =  Cu2S. 

The  formation  of  sulphides  by  the  action  of  hydrogen 
sulphide  upon  solutions  of  salts  was  discussed  at  some 
length  under  Hydrogen  Sulphide  (which  see).  The  ex- 
tensive use  made  of  this  reaction  in  chemical  analysis 
was  also  referred  to. 

The  action  of  soluble  sulphides  or  solutions  of  salts  is 
in  general  the  same  as  that  of  Irydrogen  sulphide,  but  in 
some  cases,  in  which  the  former  will  not  act,  the  latter 
will.  Thus,  hydrogen  sulphide  will  not  precipitate  iron 
sulphide  from  a  solution  of  an  iron  salt,  because  iron 
sulphide  is  easily  acted  upon  by  dilute  acids.  Thus, 
when  hydrogen  sulphide  is  passed  into  a  solution  of 
ferrous  chloride,  it  naturally  tends  to  form  the  sul- 
phide FeS: 

FeCl2  +  H2S  =  FeS  +  2HC1. 


SULPHIDES. 

But  ferrous  sulphide,  FeS,  is  acted  upon  by  dilute  hy« 
drochloric  acid,  and  is  converted  by  it  into  ferrous  chlo- 
ride and  hydrogen  sulphide : 

FeS  +  2HC1  =  FeCl2  +  H2S. 

It  is  therefore  obvious  that  the  first  reaction  cannot  take 
place. 

If,  however,  a  soluble  sulphide,  as  sodium  or  ammonr 
ium  sulphide,  is  added  to  a  solution  of  an  iron  salt,  iron 
sulphide  is  precipitated,  as  in  this  case  no  free  acid  is 
formed.  Thus,  when  ferrous  chloride  and  ammonium 
sulphide  are  brought  together  the  reaction  takes  place 
as  represented  in  the  equation 

FeCl2  +  (NH4)2S  =  FeS  +  2NH4C1. 

In  ammonium  chloride  the  ferrous  sulphide  is  not 
soluble. 

The  formation  of  a  sulphide  by  reduction  of  a  sul- 
phate is  illustrated  by  the  formation  of  barium  sulphide 
by  heating  a  mixture  of  barium  sulphate  and  charcoal : 

BaSO4  +  4C  =  BaS  +  4CO  ; 

and  by  the  formation  of  copper  sulphide  by  heating 
copper  sulphate  in  a  current  of  hydrogen : 

CuS04  +  4H2  =  CuS  +  4H20. 

The  sulphides  of  the  alkali  metals  are  soluble  in  water. 
Those  of  the  other  metals  are  insoluble.  It  should  be 
remarked,  however,  that  aluminium  and  chromium  do 
not  form  sulphides,  or,  at  least,  if  they  do,  the  compounds 
are  decomposed  by  water  into  hydroxides  and  hydrogen 
sulphide.  Barium  sulphide  is  decomposed  by  water, 
and  probably  magnesium  sulphide  also. 

The  sulphides  are  stable  when  heated  without  access 
of  air ;  but  if  heated  in  the  air  they  are  converted  into 
oxides  of  the  metals  and  sulplmr  dioxide,  or,  in  some 
cases,  they  take  up  oxygen  and  are  converted  into  sul- 
phates. The  conversion  of  sulphides  into  oxides  and 
sulphur  dioxide  by  heating  in  contact  with  the  air  has 


474  INORGANIC  CHEMISTRY. 

been  repeatedly  referred  to.  The  process  is  carried  on 
on  the  large  scale  in  the  preparation  of  iron  ores  for  re- 
duction, and  is  called  roasting.  The  conversion  of  a 
sulphide  into  a  sulphate  by  heating  is  a  simple  process 
of  oxidation.  Copper  sulphide  is  converted  into  the 
sulphate  when  heated  for  some  time  : 

CuS  +  4O  =  CuS04. 

This  is  the  reverse  of  the  reaction  mentioned  by  which  a 
sulphate  is  converted  into  a  sulphide  by  reduction. 

Some  sulphides,  as  those  of  sodium,  potassium,  and 
ammonium,  take  up  sulphur  in  much  the  same  way  that 
they  take  up  oxygen,  and  form  the  polysulphides.  The 
two  reactions  appear  to  be  entirely  analogous : 

K2S  +  40  =  K2S04 ; 

K2S  +  48  =  K2SS4,  or  K2S6. 

Hydrosulphid.es. — The  hydrosulphides  bear  the  same 
relation  to  the  sulphides  that  the  hydroxides  bear  to  the 
oxides.  They  are  not,  however,  as  numerous  nor  as 
easily  obtained  as  the  hydroxides.  When  a  hydrosul- 
phide,  as,  for  example,  potassium  hydrosulphide,  KSH, 
is  added  to  a  salt  containing  a  metal  whose  sulphide 
is  insoluble,  the  sulphide,  and  not  the  hydrosulphide, 
is  precipitated.  Thus,  copper  sulphate  and  potassium 
hydrosulphide  give  copper  sulphide  : 

CuSO4  +  2KSH  =  CuS  +  H2S  +  K2SO4. 

If  the  reaction  took  place  in  the  same  way  that  it  does  with 
the  hydroxide,  the  product  would  be  copper  hydrosul- 
phide : 

CuSO4  +  2KSH  =  Cu(SH)2  +  K2SO4. 

If  this  is  formed  it  certainly  breaks  down  into  copper 
sulphide  and  hydrogen  sulphide,  in  the  same  way  that 
copper  hydroxide  breaks  down  into  copper  oxide  and 
water,  only  more  easily  : 

Cu(SH)2  =  CuS  +  H2S ; 
Cu(OH)2  =  CuO  +  H20. 


SULPHO-SALTS.  475 

The  only  hydrosulphides  known  are  derived  from  the 
members  of  the  potassium  and  calcium  groups,  and  these 
are  soluble.  They  are  formed  by  saturating  solutions  of 
the  corresponding  hydroxides  with  hydrogen  sulphide. 
Potassium  hydrosulphide  is  formed  thus  : 

KOH  +  H2S  =  KSH  +  H2O. 
Ammonium  hydrosulphide  is  formed  thus : 

NH4OH  +  H2S  =  NH4SH  +  H2O. 

It  also  appears  probable  that  whenever  a  sulphide  is 
dissolved  in  water  it  is  converted  into  a  hydrosulphide 
and  a  hydroxide.  Thus  it  seems  to  be  true  that  potas- 
sium sulphide  is  converted  into  the  hydrosulphide  and 
hydroxide : 

K2S  +  H20  =  KSH  +  KOH. 

Sulpho-salts. — The  relation  of  the  sulpho-salts  to  the 
sulphides  has  already  been  explained.  It  is  like  that  of 
the  ordinary  oxygen  salts  to  the  oxides,  and  that  of  the 
chloro-salts,  or  double  chlorides,  to  the  chlorides.  They 
are  formed  by  dissolving  the  sulphides  of  certain  metals, 
particularly  tin,  arsenic,  and  antimony,  in  the  sulphides 
of  the  members  of  the  potassium  group  : 

As2S3  +  3K2S  =  2K3AsS3 ; 
As2S5  +  3K2S  =  2K3AsS4 ; 
SnS2  +  K2S  =  K2SnS3,  etc. 

These  sulpho-salts  are  decomposed  by  the  ordinary  acids, 
the  insoluble  sulphides  being  precipitated  thus  : 

2K3AsS3  +  6HC1  =  As2S3  +  6KC1  +  3H2S  ; 
2K3AsS4  +  6HC1  =  As2S5  +  6KC1  +  3H2S. 

Nitrates. — The  nitrates  are  formed  by  dissolving  the 
metals  in  nitric  acid,  and  by  treating  oxides,  hydroxides, 
carbonates,  and  some  other  easily  decomposed  salts  with 
nitric  acid.  The  action  of  nitric  acid  upon  metals  was 
discussed  under  the  head  of  Nitric  Acid  (which  see).  It 
was  pointed  out  that  the  acid  gives  up  a  part  of  its 


476  INORGANIC  CHEMISTRY. 

oxygen  to  the  metal  and  forms  different  oxides,  accord- 
ing to  the  conditions.  Thus,  when  the  acid  acts  upon 
copper  the  main  product  of  the  reduction  is  nitric 
oxide,  but  by  changing  the  concentration  of  the  acid  a 
considerable  quantity  of  nitrous  oxide  is  formed.  When 
zinc  is  dissolved  in  nitric  acid  a  part  of  the  acid  is 
reduced  to  ammonia. 

The  nitrates  are  soluble  in  water,  and  all  are  decom- 
posed by  heat.  Some  of  them  when  heated  lose  only  a 
third  of  their  oxygen  and  are  reduced  to  nitrites.  This 
is  true  of  potassium  nitrate,  the  decomposition  of  which 
is  represented  by  the  equation 

KNO3  =  KNO2  +  O. 

Most  of  the  nitrates,  however,  are  decomposed  further, 
forming  oxides.  This  has  been  shown  in  the  case  of  lead 
nitrate,  which  when  heated  is  converted  into  lead  oxide, 
while  nitrogen  peroxide  and  oxygen  are  given  off: 

Pb(N03)2  =  PbO  +  2NO2  +  O. 

If  the  oxide  of  the  metal  is  decomposed  by  heat,  as  in 
the  case  of  mercury,  of  course  the  product  will  be  the 
metal. 

Chlorates. — These  salts,  except  potassium  chlorate,  are 
not  commonly  met  with.  Potassium  chlorate  is  manu- 
factured in  large  quantity,  and  the  other  chlorates  are 
generally  made  from  it.  The  chlorates  are  soluble  in 
water,  and  are  decomposed  by  heat  more  easily  than  the 
nitrates  are.  They  are  first  converted  into  perchlorates, 
and  these  are  further  decomposed  by  higher  heat  into 
chlorides  and  oxygen. 

The  hypocJdorites  are  formed  by  treating  some  of  the 
metallic  hydroxides  in  dilute  solution  with  chlorine. 
This  has  been  illustrated  in  the  formation  of  "  bleaching 
powder,"  which  contains  calcium  hypochlorite  or  a  com- 
pound closely  related  to  it.  The  hypochlorites,  like  the 
chlorates,  are  decomposed  by  heat. 

Sulphates. — The  general  relations  of  the  sulphates  to 
sulphuric  acid  were  treated  of  under  Sulphuric  Acid 
(which  see).  Some  of  these  salts  occur  in  nature  in  large 


SULPHATES.  477 

quantity,  as  those  of  calcium  and  barium.  The  former 
is  known  as  gypsum,  the  latter  as  heavy  spar.  Sulphates 
are  made  by  treating  metals,  metallic  hydroxides  or  ox- 
ides, carbonates,  etc.,  with  sulphuric  acid ;  and  by  treat- 
ing a  solution  containing  a  metal  whose  sulphate  is 
insoluble,  with  sulphuric  acid  or  a  soluble  sulphate. 

Zinc  and  iron  give  hydrogen  and  a  sulphate  when 
treated  with  sulphuric  acid  : 

Zn  +  H2S04  =  ZnSO4  +  H2 ; 
Fe  +  H2S04  =  FeS04  +  H2. 

This  kind  of  action  takes  place  whenever  a  metal  is  dis- 
solved in  sulphuric  acid  at  the  ordinary  temperature.  If, 
however,  the  temperature  is  raised  the  displaced  hydro- 
gen acts  upon  some  of  the  sulphuric  acid,  or  the  metal 
extracts  some  of  the  oxygen  of  the  acid,  reducing  it 
partly  to  sulphurous  acid,  when  sulphur  dioxide  is  given 
off.  This  happens  in  the  case  of  copper,  as  has  been 
pointed  out.  It  may  be  represented  either  by  these 
•equations  : 

Cu        +  H2SO4  =  CuS04  +  H2 ; 
H3S04  +  2H       =  SOa       +  2H2O ; 

or  by  these : 

Cu     +  H2SO4  =  CuO     +  SO2  +  HaO ; 
CuO  +  H2SO4  =  CuSO4  +  H3O. 

The  action  of  sulphuric  acid  on  metallic  hydroxides 
has  been  fully  described. 

Most  sulphates  are  soluble  in  water.  The  sulphates 
of  barium,  strontium-,  and  lead  are  insoluble  in  water, 
and  the  sulphate  of  calcium  is  difficultly  soluble.  There- 
fore, if  sulphuric  acid  or  a  soluble  sulphate  is  added  to 
a  solution  containing  either  of  the  metals,  barium,  stron- 
tium, or  lead,  a  precipitate  is  formed.  A  precipitate 
is  also  formed  when  a  concentrated  solution  of  a  calcium 
salt  is  treated  in  the  same  way. 


478  INORGANIC  CHEMISTRY. 

When  heated  with  charcoal  in  the  reducing  flame  of 
the  blow-pipe,  sulphates  are  reduced  to  sulphides : 

K2SO4  +  40  =  K2S  +  4CO,  or 
K2SO4  +  2C  =  K3S  +  2C02. 

Sulphites  are  made  from  sodium  or  potassium  sul- 
phite, which  are  made  by  treating  sodium  or  potassium 
hydroxide  in  solution  with  sulphur  dioxide : 

2NaOH  +  SO2  =  Na2SO3  +  H2O. 

All  sulphites  are  decomposed  by  the  common  acids, 
sulphur  dioxide  being  given  off : 

Na2SO3  +  H2SO4  =  Na2SO4  +  H2O  +  SO,. 

The  sulphites  are  changed  to  sulphates  by  oxidation. 
Thus,  sodium  sulphite  is  changed  to  the  sulphate  when 
its  solution  is  allowed  to  stand  in  contact  with  the  air : 

Na2S03  +  O  =  Na2SO4. 

The  sulphites,  like  the  sulphates,  are  reduced  to  sul- 
phides. 

Carbonates. — Many  carbonates  are  found  in  nature, 
some  of  them  in  great  abundance  and  widely  distrib- 
uted. The  principal  one  is  calcium  carbonate.  They 
are  made  by  passing  carbon  dioxide  into  solutions  of  hy- 
droxides, and  by  adding  soluble  carbonates  to  solutions 
of  salts  containing  metals  whose  carbonates  are  insolu- 
ble. 

The  formation  of  carbonates  by  the  action  of  carbon 
dioxide  on  a  solution  of  hydroxide  is  illustrated  in  the 
case  of  potassium  hydroxide  : 

2KOH  +  CO2  =  K2C03  +  H2O. 

The  formation  of  calcium  carbonate  takes  place  in  the 
same  way,  but  the  carbonate  formed  is  insoluble  : 

Ca(OH)2  +  C02  =  CaCO3  +  H,0. 


PHOSPHATES.  479 

If  in  either  case  the  action  is  continued,  the  normal 
carbonate  first  formed  is  converted  into  acid  carbonate : 

K2CO3  +  CO,  +  H2O  =  2KHCO3 ; 
CaCO3  +  CO2  +  H2O  =  Ca  )  ,pn  x 

H2  f  v^Vt/*1 

All  normal  carbonates  except  those  of  the  members  of 
the  potassium  family  are  insoluble,  and  are  decomposed 
by  heat  into  carbon  dioxide  and  the  oxide  of  the  metal. 
The  decomposition  of  calcium  carbonate  into  lime  and 
carbon  dioxide  is  an  example : 

CaCO3  =  CaO  +  CO2. 

When  a  soluble  carbonate  is  added  to  a  solution  of  a 
calcium,  barium,  or  strontium  salt,  the  corresponding 
insoluble  carbonates  are  precipitated.  When  a  magne- 
sium salt  is  treated  with  a  soluble  carbonate,  however,  a 
basic  carbonate  is  precipitated  : 

4MgSO4  +  3Na2CO3  +  2H2O  =  Mg4(OH)2(CO3)3  +  3Na2SO4  +  H2SO4. 

This  salt,  which  at  first  sight  appears  to  be  quite  com- 
plicated, is  in  all  probability  derived  from  three  mole- 
cules of  carbonic  acid  and  four  of  magnesium  hydroxide, 
as  represented  in  the  formula  on  page  388.  Many  other 
metals  give  basic  carbonates  under  the  same  conditions. 
Further,  some  of  the  metals,  like  aluminium,  chromium, 
and  tin  do  not  form  salts  with  carbonic  acid.  If,  there- 
fore, salts  of  these  metals  are  treated  with  soluble  car- 
bonates, the  oxides  or  hydroxides  are  thrown  down,  and 
not  the  carbonates. 

Phosphates. — Calcium  phosphate  is  very  abundant  in 
nature,  and  a  few  other  phosphates  are  also  found.  The 
methods  used  for  making  phosphates  are  the  same  as 
those  used  in  making  salts  in  general. 

The  normal  phosphates  of  all  the  metals  except  the 
members  of  the  potassium  family  are  insoluble  in  water. 
The  normal  phosphates,  as  a  rule,  are  not  changed  by 
heat.  The  secondary  phosphates,  such  as  secondary 


480  INORGANIC  CHEMISTRY. 

sodium  phosphate,  HNaaPO4,  lose  water  when  heated,, 
and  yield  pyrophosphates  : 

2HNaaPO4  =  Na4P2O7  +  H2O. 

Sodium 
pyrophosphate. 

Those  phosphates  in  which  only  one  third  of  the  hy- 
drogen is  replaced  by  metal — as,  for  example,  primary 
sodium  phosphate,  HaNaPO4 — lose  water  when  heated, 
and  yield  metaphosphates  : 

HaNaPO4  =  NaPO3  +  H2O. 

Sodium 
metaphosphate 

Neither  the  pyrophosphates  nor  the  metaphosphates  are 
changed  by  heat. 

Silicates. — The  silicates,  as  has  been  stated,  are  very 
widely  distributed  in  nature.  Those  which  are  most 
abundant  are  the  feldspars  and  their  decomposition-pro- 
ducts. The  principal  feldspar  is  a  complex  silicate  of 
aluminium  and  potassium,  of  the  formula  KAlSi3O8,  de- 
rived from  the  polysilicic  acid  H4Si3O8,  which  is  formed 
from  three  molecules  of  normal  silicic  acid  by  the  loss 
of  four  molecules  of  water : 

3Si(OH)4  =H4Si3O8  +  4HSO. 

Silicates  can  be  made  by  heating  together,  at  a  high 
temperature,  silicon  dioxide,  in  the  form  of  sand,  and 
basic  oxides  or  carbonates  : 

CaO      +  SiOa  =  CaSi03 ; 
NaaCO3  +  SiO2  =  NaaSi03  +  C02. 

Only  the  silicates  of  the  members  of  the  potassium 
group  are  soluble  in  water.  When  these  are  treated  in 
solution  with  dilute  acids,  they  are  decomposed,  as  has 
been  explained  under  Silicic  Acid  (which  see). 

Some  silicates,  which  are  insoluble  in  water,  are  decom- 
posed by  the  ordinary  acids,  such  as  sulphuric  and  hy- 
drochloric acids,  the  silicic  acid  separating  as  a  difficultly 
soluble  substance,  which,  if  dried  on  the  water-bath,  be- 
comes insoluble. 


SILICATES.  481 

Many  silicates,  which  are  not  acted  upon  by  strong 
acids,  are  decomposed  by  fusion  with  sodium  or  potassi- 
um carbonate,  when  the  silicate  of  potassium  or  sodium 
and  the  oxide  of  the  metal  contained  in  the  silicate  are 
formed.  Silicates  which  are  not  decomposed  in  either 
of  the  ways  mentioned,  yield  to  hydrofluoric  acid.  The 
action  consists  in  the  formation  of  the  gas,  silicon  tetra- 
fluoride,  SiF4,  and  the  fluorides  of  the  metals  present. 
Thus,  the  reaction  in  the  case  of  feldspar  takes  place  in 
accordance  with  the  equation, 

KAlSi3O8  +  16HF  =  KF  +  A1F3  +  3SiF4  +  8H2O. 

The  silicon  fluoride  is  given  off  as  a  gas,  and  the  flu- 
orides formed  are  soluble  in  water.  Hence,  hydrofluoric 
acid  is  said  to  dissolve  the  silicates. 


CHAPTER  XXV0 

ELEMENTS  OF  FAMILY  I,    GROUP  A: 

THE  ALKALI  METALS  :—  LITHIUM—  SODIUM—  POTASSIUM- 
RUBIDIUM—  CJESIUM—  AMMONIUM. 

General.  —  The  elements  of  this  group  which  are  most 
abundant  in  nature  are  sodium  and  potassium.  While 
lithium  occurs  in  considerable  quantity,  the  two  remain- 
ing elements,  rubidium  and  caesium,  have  been  found  in 
only  very  small  quantities.  ,  They  are  all  strongly  basic, 
their  hydroxides  being  the  strongest  bases  known.  They 
form  well-  characterized  salts  with  all  acids,  and  as  a 
rule  their  salts  are  very  stable.  In  all  their  compounds 
they  act  as  univalent  elements,  except  in  those  which 
they  form  with  hydrogen,  and  in  their  peroxides  ;  in  the 
latter  they  appear  to  be  bivalent.  Leaving  these  com- 
pounds out  of  consideration  the  general  formulas  of 
some  of  the  other  principal  compounds  are  as  follows  : 

MCI,  M2O,  M2S,  M(OH),  M(SH),  MNO3,  M2SO4,  etc. 

The  valence  of  the  members  of  the  group  towards  other 
elements  is,  in  general,  constant. 

The  relations  between  the  atomic  weights  are  interest- 
ing. That  of  sodium,  22.82,  is  very  nearly  half  the  sum  of 
those  of  lithium,  6.97,  and  potassium,  38.82.  We  have 

-  ("7      88.* 


i 

A 

So,  also,  that  of  rubidium,  84.78,  is  approximately  half 
the  sum  of  those  of  potassium,  38.82,  and  caesium,  131.89. 


38.82  + 131.89 

~ =  oD.OO. 

a 


(482) 


POTASSIUM.  483 

The  specific  gravity  of  these  elements  increases  with 
the  atomic  weight;  and  their  melting-points  become 
lower  as  the  atomic  weights  become  higher. 

At.  Wt.  Sp.  Grav.  M.  P. 

Lithium,    .     .    .        6.97  0.594  180°. 

Sodium,    .     .    .       22.82  0.972  95.6° 

Potassium,     .     .       38.82  0.865  62.5° 

Kubidium,     .     .      84.78  1.52  38.5° 

Cesium,    .     .     .     131.89  ?  ? 

The  regularity  is  complete  in  the  case  of  the  melting- 
points,  but  as  regards  the  specific  gravities  sodium  is  an 
exception  to  the  rule.  As  sodium  and  potassium  and 
their  compounds  are  much  more  commonly  met  with 
than  the  other  members  of  the  group,  these  will  form 
the  chief  subject  of  consideration  in  this  chapter. 

POTASSIUM,  K  (At.  Wt.  38.82). 

Occurrence. — Potassium  is  a  constituent  of  many  min- 
erals, particularly  of  feldspar,  the  common  variety  of 
which,  as  has  already  been  explained,  is  a  complex  sili- 
cate of  aluminium  and  potassium.  It  is  found  also  in 
combination  with  chlorine  as  carnallite  and  sylvite ;  with 
sulphuric  acid  and  aluminium,  as  alum  ;  with  nitric  acid, 
as  saltpeter  or  potassium  nitrate ;  and  in  other  forms. 
The  natural  decomposition  of  minerals  containing  potas- 
sium gives  rise  to  the  presence  of  this  metal  in  various 
forms  of  combination  everywhere  in  the  soil.  It  is  taken 
up  by  the  plants ;  and,  when  vegetable  material  is  burned, 
the  potassium  remains  behind,  chiefly  as  potassium  car- 
bonate. When  wood-ash  is  treated  with  water  the 
potassium  carbonate  is  dissolved,  and  it  can  be  obtained 
in  an  impure  state  by  evaporating  the  solution.  The 
substance  thus  obtained  is  called  potash.  In  the  juice  of 
the  grape  there  is  contained  a  salt  of  potassium,  mono- 
potassium  tartrate,  which  is  deposited  in  large  quantity 
from  wine.  This  is  commonly  called  "crude  tartar." 

Preparation. — Potassium  was  first  prepared  by  Davy 
in  the  year  1807,  by  the  action  of  a  powerful  electric  cur- 
rent on  potassium  hydroxide.  It  is  now  prepared  by 


484  INORGANIC  CHEMISTRY. 

heating  to  a  high  temperature  a  mixture  of  potassium 
carbonate  and  carbon : 

KaCO3  +  20  =  2K  +  3CO. 

Such  a  mixture  is  best  obtained  by  heating  in  a  closed 
vessel  ordinary  mono-potassium  tartrate  obtained  from 
wine.  This  contains  some  calcium  tartrate.  When  the 
whole  is  heated  decomposition  takes  place,  and  there  is 
left  behind  an  intimate  mixture  of  potassium  carbonate, 
calcium  carbonate,  and  charcoal.  This  mixture  is  placed 
in  a  wrought-iron  retort  which  is  connected  with  a  closed 
flat  receiver  of  sheet-iron.  The  retort  is  then  heated  to 
a  high  temperature.  The  metal  distils  over  into  the 
closed  receiver,  and  at  the  end  of  the  operation  the  re- 
ceiver is  placed  under  petroleum  to  protect  it  from  the 
action  of  the  air.  The  metal  obtained  in  this  way  is 
not  pure.  It  can  be  partly  purified  by  melting  it  under 
petroleum  and  pressing  it  through  a  linen  bag.  It  can 
also  be  purified,  and  more  completely,  by  distilling  it 
from  a  wrought-iron  retort. 

Properties. — Potassium  is  a  light  substance,  which 
floats  on  water.  Its  freshly  cut  surface  has  a  bright 
metallic  lustre,  almost  white ;  it  acts  energetically  upon 
water,  causing  the  evolution  of  hydrogen,  which,  together 
with  some  of  the  potassium,  burns,  while  potassium  hy- 
droxide is  formed  at  the  same  time.  This  reaction  has 
been  studied  in  connection  with  hydrogen.  In  conse- 
quence of  its  action  upon  water,  potassium  cannot  be 
kept  in  the  air.  It  is  kept  under  some  oil,  as  petroleum, 
upon  which  it  does  not  act.  In  an  atmosphere  upon 
which  it  does  not  act,  as,  for  example,  hydrogen,  it  can 
be  distilled.  Its  vapor  is  green.  Its  specific  gravity  is 
0.865;  its  melting-point  62.5°.  It  combines  with  chlo- 
rine and  bromine  with  great  energy,  and  has  the  power 
to  extract  chlorine  from  its  compounds.  It  can,  there- 
fore, be  used  for  the  purpose  of  isolating  some  elements, 
as,  for  example,  magnesium  and  aluminium,  whose  oxy- 
gen compounds  cannot  be  reduced  by  the  ordinary 
methods.  As,  however,  sodium  is  generally  used  for 
this  purpose  instead  of  potassium -,  on  account  of  its  lower 


POTASSIUM  SALTS.  485 

price,  the  action  will  be  referred  to  more  at  length  under 
Sodium.  Although  the  metal  is  converted  into  vapor, 
no  reliable  determination  of  the  specific  gravity  of  the 
vapor  has  been  made,  for  the  reason  that  the  vessels 
which  have  been  used  for  the  purpose  have  always 
been  acted  upon,  and  the  results  thus  vitiated. 

Potassium  Hydride,  K2H. — This  compound  is  formed 
by  heating  potassium  in  an  atmosphere  of  hydrogen  at 
about  300°.  It  is  a  silver-white  mass  with  a  metallic 
lustre.  It  takes  fire  in  the  air.  When  heated,  it  begins 
to  dissociate  at  200°. 

[  Fluoride,  KF 

Potassium]  gjEdd?;  f&  —  Of   these   salts  the  only 

L  Iodide,  KI 

one  which  occurs  in  nature  in  quantity  is  the  chloride. 
This  is  found  in  the  great  salt  deposits  at  Stassfurt, 
Germany,  and  in  some  other  localities  in  the  form  of  the 
mineral  sylvite,  which  is  more  or  less  impure  potassium 
chloride.  It  is  also  found  in  the  form  of  a  compound 
containing  magnesium,  potassium,  and  chlorine,  of  the 
formula  MgCl2.KCl  +  6H2O,  or  KMgCl3  +  6H2O,  known 
as  carnallite. 

The  other  salts  of  the  group  are  made  by  the  general 
methods  for  making  salts,  that  is,  by  neutralizing  the 
acids  with  the  hydroxide  or  carbonate  of  potassium.  It 
is,  however,  easier  to  make  the  iodide  by  other  methods, 
and  as  there  is  a  large  demand  for  this  salt  for  use  in 
medicine  and  in  the  art  of  photography,  several  methods 
have  been  devised  for  its  preparation.  Of  these,  two 
may  serve  as  examples  :  (1)  The  first  consists  in  treating 
a  solution  of  potassium  hydroxide  with  iodine  until  it 
begins  to  show  a  permanent  yellow  color,  which  is  an 
indication  that  no  more  iodine  will  be  taken  up.  The 
action  is  the  same  as  that  which  takes  place  when  chlo- 
rine acts  upon  warm  concentrated  caustic  potash.  Both 
the  iodide  and  iodate  are  formed  : 

6KOH  +  61  =  SKI  +  KIO3  +  3H2O. 

By  evaporating  the  water  and  heating  the  residue  with 
very  finely  powdered  charcoal,  the  iodate  is  decomposed 


486  INORGANIC  CHEMISTRY. 

into  iodide  and  oxygen.  The  reduction  of  the  iodate 
takes  place  in  accordance  with  the  equation : 

2KIO3  +  30  =  2KI  +  3CO2. 

(2)  Another  method  employed  in  the  preparation  of 
potassium  iodide  consists  in  treating  iron  filings  under 
water  with  iodine.  Both  the  iron  and  iodine  dissolve, 
forming  ferrous  iodide,  FeI2.  If  to  the  solution  of  this 
compound  half  as  much  iodine  is  added  as  has  already 
been  used  in  its  preparation,  ferroso-ferric  iodide,  Fe3I8, 
is  formed  and  remains  in  solution.  By  adding  a  solution 
of  potassium  carbonate  to  this,  reaction  takes  place  as 
represented  in  the  equation  : 

Fe3I.  +  4K2C03  +  4H20  =  SKI  +  Fe3(OH)8  +  4CO2t 

The  hydroxide  of  iron  is  insoluble,  and  can  be  removed 
by  filtration. 

The  fact  that  the  specific  gravity  of  hydrofluoric  acid 
at  a  low  temperature  corresponds  to  the  formula  H2F2 
makes  it  not  improbable  that  potassium  fluoride  has 
the  formula  K2F2.  This  appears  still  more  probable 
from  the  fact  that  there  is  an  acid  potassium  fluoride  of 
the  formula  KHF2,  or  KF  +  HF.  Similar  acid  salts 
have  not  been  obtained  from  the  other  acids  of  the 
group. 

Properties. — All  these  salts  are  soluble,  and  crystallize 
well  in  cubes.  The  fluoride  is  the  most  easily  soluble 
in  water.  If  deposited  from  a  water  solution  at  the  or- 
dinary temperature  the  crystals  contain  two  molecules 
of  water  of  crystallization,  and  are  deliquescent.  The 
iodide  is  soluble  in  0.7  parts  of  water  at  the  ordinary 
temperature,  and  is  also  soluble  in  alcohol  (40  parts). 
The  bromide  requires  about  1^  parts  of  water  for  solu- 
tion at  the  ordinary  temperature,  and  is  but  slightly 
soluble  in  alcohol.  The  chloride  is  soluble  in  3  parts  of 
water  at  the  ordinary  temperature,  and  is  insoluble  in 
alcohol.  All  are  decomposed  by  sulphuric  acid.  The 
fluoride  gives  hydrofluoric  acid ;  the  chloride  gives  hy- 
drochloric acid.  The  bromide  gives  hydrobromic  acid, 
which  acts  upon  the  sulphuric  acid,  giving  sulphur  di- 


DOUBLE  HALOGEN  SALTS. 


48: 


oxide  and  free  bromine  (see  Hydrobromic  Acid).  The 
action  in  the  case  of  the  iodide  is  more  complicated  for 
the  reason  that  hydriodic  acid  is  less  stable  than  hydro- 
bromic  acid,  and,  as  it  gives  up  hydrogen  very  easily,  it 
causes  deeper-seated  decomposition  of  the  sulphuric 
acid  (see  Hydriodic  Acid).  Potassium  iodide  in  solution 
takes  up  iodine  readily,  and  a  compound  of  the  formula 
KI3  can  be  isolated  from  a  very  concentrated  solution. 
No  similar  compounds  of  chlorine,  bromine,  and  fluorine 
are  known. 

All  the  salts  of  this  group  combine  readily  with  the 
fluorides,  chlorides,  bromides,  and  iodides  of  the  metallic 
elements  in  general,  forming  salts,  of  which  the  double 
fluorides  and  double  chlorides  are  examples.  The  re- 
lations between  these  salts  and  the  ordinary  oxygen 
salts  have  already  been  discussed  to  a  sufficient  extent 
(see  pp.  465  and  466).  Those  containing  fluorine  have 
been  studied  most  fully.  Good  examples  are  the  fol- 
lowing : 


F 
F 

As^jF     ,orAsFB.KF; 
F 
F2K 


rr 

F 

F     ,  or  AsF6.2KF ; 
F2K 
[F2K 


F 

F  (F 

F     ,orSbF5.2KF;   B^F     ,orBF3.KF; 

F2K  F2K 

F2K 


Applications.— Potassium  chloride  is  extensively  used 
for  the  purpose  of  making  other  potassium  salts,  as,  for 
example,  the  nitrate  and  carbonate  ;  the  bromide  is  used 
in  medicine  ;  the  iodide,  as  stated  above,  is  used  in  medi- 
cine and  in  photography. 


488  INORGANIC  CHEMISTRY. 

Potassium  Hydroxide,  KOH. — This  well-known  sub- 
stance, commonly  called  caustic  potash,  is  prepared  by 
treating  potassium  carbonate  in  solution  with  calcium 
hydroxide  in  a  silver  or  iron  vessel.  The  reaction  is 
based  upon  the  fact  that  calcium  carbonate  is  insoluble, 
and  that  potassium  carbonate  and  calcium  hydroxide 
are  soluble  : 

K2CO3  +  Ca(OH)2  =  2KOH  +  CaCO3. 

After  enough  lime  has  been  added,  it  is  found  that  a 
little  of  the  liquid  taken  out  of  the  vessel  gives  no  carbon 
dioxide  when  treated  with  acids.  When  this  point  is 
reached  the  liquid  is  drawn  off  from  the  deposit  of  cal- 
cium carbonate  by  means  of  a  siphon.  In  the  prepara- 
tion on  the  large  scale  this  is  then  evaporated  down  in  a 
bright  wrought-iron  vessel  until  it  has  the  specific  gravity 
1.16.  If  the  evaporation  is  carried  farther  the  liquid 
acts  upon  the  iron.  Concentration  beyond  this  point 
must  be  carried  on  in  silver  vessels,  upon  which  potas- 
sium hydroxide  does  not  act.  Finally,  a  liquid  is  ob- 
tained which  on  cooling  completely  solidifies.  While  in 
the  molten  condition  it  is  generally  poured  into  moulds  of 
cast-iron  or  of  brass,  plated  with  silver,  in  which  it  solidi- 
fies in  the  form  of  the  thin  sticks  found  in  the  market. 
This  substance  is  generally  not  pure.  It  always  con- 
tains some  carbonate  formed  by  the  action  of  the  carbon 
dioxide  of  the  air,  and  other  substances  are  also  present 
in  small  quantity.  It  can  be  purified  by  dissolving  it 
in  alcohol,  in  which  the  impurities  are  insoluble.  The 
alcoholic  solution  of  the  hydroxide  is  poured  off  after  a 
time  and  evaporated  to  dryness  in  a  silver  vessel.  The 
liquid  becomes  colored  in  consequence  of  a  partial  de- 
composition of  the  alcohol,  but  on  melting  the  residue 
the  color  disappears,  as  the  substances  formed  from  the 
alcohol  are  thus  destroyed.  This  product  is  known  as 
"  caustic  potash  by  alcohol."  Pure  potassium  hydroxide 
in  solution  is  easily  obtained  by  the  action  of  potassium 
upon  distilled  water. 

Potassium  hydroxide  is  a  white  brittle  substance.  In 
contact  with  the  air  it  deliquesces,  and  absorbs  carbon 


POTASSIUM  OXIDE.  489 

dioxide,  being  completely  transformed  into  potassium 
carbonate.  It  is  the  strongest  of  the  bases.  It  decom- 
poses the  salts  of  all  other  bases,  even  of  those  which, 
like  sodium  and  lithium  hydroxides,  are  soluble  in  water. 

Animal  substances  like  the  skin  are  disintegrated  by 
the  hydroxide.  It  has  a  caustic  action.  It  is  interest- 
ing to  observe  that  the  strongest  bases,  like  the  strongest 
acids,  exert  this  kind  of  influence  on  the  complex  organic 
compounds  which  go  to  make  up  the  tissues  of  animals. 
The  action  is  not  by  any  means  always  of  the  same  kind, 
and  all  that  can  be  said  in  regard  to  it,  of  a  general 
character,  is  that  it  tends  to  break  down  the  complex 
substances  to  simpler  ones.  In  the  molten  condition 
the  hydroxide  acts  as  an  oxidizing  agent.  Hydrogen  is 
given  up  from  it,  and  substances  of  acid  character  are 
formed  with  which  the  potassium  combines,  forming 
salts. 

Instead  of  potassium  hydroxide,  the  corresponding 
sodium  compound  is  used  wherever  this  is  possible,  as 
the  latter  is  cheaper.  The  chief  application  of  the 
potassium  compound  outside  of  the  laboratory  is  for 
making  soft-soap.  For  this  purpose  fats  are  boiled  with 
a  solution  of  potassium  hydroxide  or  carbonate. 

Potassium  Oxide,  K2O. — This  compound  can  be  made 
by  burning  potassium  in  the  air,  and  heating  the  residue 
to  a  high  temperature.  It  is  also  formed  by  melting 
potassium  hydroxide  and  metallic  potassium  together  : 

2K  +  2KOH  =  2K2O  +  Ha. 

With  water  it  forms  the  hydroxide,  with  a  marked  evo- 
lution of  heat : 

K20  +  H20  =  2KOH. 

Potassium  also  forms  other  oxides  of  which  the  peroxidt 
of  the  formula  K2O4  is  the  best  studied.  This  peroxide 
is  the  final  product  of  the  combustion  of  potassium  in 
the  air  or  in  oxygen.  At  a  high  temperature  it  breaks 
down  into  potassium  oxide,  K2O,  and  oxygen.  It  also 
gives  up  its  oxygen  very  readily  to  substances  which  are 


490  INORGANIC  CHEMISTRY. 

capable  of  oxidation,  acting  so  energetically  upon  some 
as  to  cause  evolution  of  light. 

Potassium  Hydrosulphide,  KSH,  is  analogous  to  potas- 
sium hydroxide.  Just  as  the  latter  is  made  by  the  action 
of  potassium  on  water,  so  the  former  can  be  made  by  the 
action  of  potassium  on  hydrogen  sulphide  : 

K2  +  2H2S  =  2KSH  +  H2. 

It  is,  however,  obtained  most  readily  by  the  action  of 
hydrogen  sulphide  on  a  solution  of  potassium  hydroxide  : 

KOH  +  H2S  =  KSH  +  H20. 

When  exposed  to  the  action  of  the  air  it  is  oxidized,  and 
becomes  colored  in  consequence  of  the  formation  of  the 
disulphide.  The  action  takes  place  as  represented  thus  : 

2KSH  +  O  =  K2S2  +  HaO. 

Potassium  Sulphide,  K2S,  is  made  by  the  reduction  of 
potassium  sulphate  either  by  means  of  hydrogen  or  car- 
bon. It  is  thought  by  some  to  be  present  in  a  solution 
prepared  by  saturating  a  given  quantity  of  potassium 
hydroxide  with  hydrogen  sulphide,  and  then  adding  the 
same  quantity  of  potassium  hydroxide  to  the  product. 
The  formation  is  supposed  to  take  place  as  represented 
in  the  equations 


KSH  +  KOH  =  K2S    +  H20. 

This  is  the  action  which  we  should  expect,  as  hydrogen 
sulphide  acts  like  an  acid,  and  with  a  strong  base  we 
should  expect  it  to  form  two  salts,  the  acid  salt  KSH 
and  the  neutral  salt  K2S.  From  thermo-chemical  investi- 
gations of  this  subject,  however,  the  conclusion  appears 
to  be  justified  that  the  salt  K2S  does  not  exist  in  solution, 
but  that  it  breaks  down  with  water,  forming  the  hydro- 
sulphide  and  hydroxide  : 

K2S  +  H2O  =  KSH  +  KOH. 

This  reaction  is  analogous  to  that  of  water  on  the  oxide  . 
K,0  +  H,O  =  2KOH. 


POLYSULPHIDES  OF  POTASSIUM.  491 

The  poly  sulphides  of  potassium  are  compounds  having 
the  composition  expressed  by  the  formulas  K2S2,  K2S3, 
K2S4,  and  K2S5.  They  are  formed  in  general  by  the 
action  of  sulphur  on  a  solution  of  the  hydrosulphide  or 
of  the  simple  sulphide.  The  disulphide  is  also  formed 
as  explained  above  by  oxidation,  when  a  solution  of  the 
hydrosulphide  is  allowed  to  stand  exposed  to  the  air. 
They  are  all  colored  substances,  which  readily  give  up 
sulphur.  If  treated  with  dilute  acids  each  one  gives  up 
sufficient  sulphur  to  reduce  it  to  the  simple  form  K2S. 
If  the  air  is  allowed  to  act  upon  them  for  a  sufficient 
length  of  time  they  all  yield  the  thiosulphate,  K2S2O3, 
the  action  taking  place  as  represented  in  the  following 
equations  : 


K,S,  +  30  =  K,S  A  +  S  ; 
K,S4  +  3O  =  K,SA  +  2S; 
K  A  +  30  =  K.S  A  +  38. 

The  fact  that  no  higher  sulphide  of  potassium  than 
the  pentasulphide  exists,  suggests  that  the  action  of 
sulphur  upon  the  monosulphide  is  analogous  to  that  of 
oxygen,  and  that  the  pentasulphide  is  analogous  to  the 
sulphate  : 

K2S  +  40  =  K2S04  ; 

K,S  +  4S  =  K2SS4,  or  K2S6. 

According  to  this  view,  the  pentasulphide  is  the  salt  of 

(mr\ 
S2S<Sjj),  or  hydrogen 

pentasulphide,  H2S5. 

The  substance  used  in  medicine  under  the  name  of 
liver  of  sulphur  or  Hepar  sulfuris  is  a  brown  mass 
formed  by  melting  together  potassium  carbonate  and 
sulphur,  and  consisting  of  polysulphides  of  potassium 
and  potassium  thiosulphate  and  sulphate.  The  chief  re- 
action involved  is  the  one  represented  in  the  equation 

3K2C03  +  8S  =  2K2S3  +  K2S2O3  +  3CO 


23  a. 


492  INORGANIC  CHEMISTRY. 

If  the  mass  is  ignited  the  thiosulphate  is  decomposed, 
forming  the  sulphate  and  pentasulphide : 

4K,S,0.  =  3K£04  +  K2S8. 

Potassium  sulphide  combines  with  the  sulphides  of 
arsenic,  antimony,  and  tin,  forming  salts  of  sulpho-acids. 
Among  the  best  known  of  these  are  the  following  : 

Potassium  Sulpharsenate,  K3AsS4  -|-  H2O.  —  This  is 
formed  by  treating  arsenic  pentasulphide  or  the  trisul- 
phide  and  sulphur  with  potassium  hydrosulphide : 

6KSH  +  As2S5  =  2K3AsS4  +  3H2S ; 
6KSH  +  As2S3  +  28  =  2K3AsS4  +  3H2S. 

It  is  also  formed  by  saturating  a  solution  of  potassium 
arsenate  with  hydrogen  sulphide. 

Potassium  sulphantimonate,  K3SbS4  +  4^H2O,  and  potas- 
sium sulpharsenite,  KAsS2  -\-  2^H2O,  are  also  easily  made. 
The  latter  is  plainly  the  analogue  of  the  metarsenite, 
KAsOa — metarsenious  acid  being  derived  from  normal 
arsenious  acid  by  elimination  of  one  molecule  of  water : 

H3AsO3  -  HAs02  +  H2O. 

Potassium  Nitrate,  KNO3. — This  salt  is  commonly  called 
saltpeter.  Its  occurrence  in  nature  has  already  been 
spoken  of  under  Nitric  Acid  (which  see).  When  refuse 
animal  matter  is  left  to  undergo  decomposition  in  the 
presence  of  bases  nitrates  are  always  the  end-products. 
They  are  consequently  found  very  widely  distributed  in 
the  soil.  In  the  East  Indies  the  potassium  nitrate 
formed  in  the  neighborhood  of  dwellings  and  stables  is 
collected,  and  sent  into  the  market.  The  process  of  nit- 
rification is  carried  on  artificially  on  the  large  scale  in 
the  so-called  "  saltpeter  plantations."  In  these,  refuse 
animal  matter  is  mixed  with  earthy  material,  wood 
ashes,  etc.,  and  piled  up.  These  piles  are  moistened 
with  the  liquid  products  from  stables.  After  the  action 
has  continued  for  two  or  three  years  the  outer  crust  is 
taken  off,  and  extracted  with  water.  The  solution  thus 
obtained  contains,  besides  potassium  nitrate,  calcium  and 
magnesium  nitrates.  It  is  treated  with  a  water-extract 


GUNPOWDER.  493 

of  wood  ashes  or  with  potassium  carbonate,  by  which  the 
calcium  and  magnesium  are  thrown  down  as  carbonates. 
Much  of  the  saltpeter  which  is  now  in  the  market  is 
made  from  Chili  saltpeter,  or  sodium  nitrate,  by  treating 
it  with  potassium  chloride,  advantage  being  taken  of  the 
fact  that  sodium  chloride  is  less  soluble  in  water  than 
potassium  nitrate.  Molecular  weights  of  sodium  nitrate 
and  potassium  chloride  are  dissolved  in  water  and  the 
solution  evaporated,  when  sodium  chloride  is  deposited : 

NaNO3  +  KC1  =  KNO3  +  NaOl. 

Potassium  nitrate  crystallizes  in  long  rhombic  prisms, 
of  a  salty  taste.  Under  some  circumstances  it  crystal- 
lizes in  rhombohedrons.  When  dissolved  in  water  it 
causes  a  lowering  of  the  temperature.  At  ordinary  tem- 
peratures 100  parts  of  water  dissolve  from  20  to  30  parts 
of  the  salt ;  at  100°,  100  parts  dissolve  247  parts. 

Applications. — Potassium  nitrate  is  used  as  an  oxidiz- 
ing agent  in  the  laboratory,  and  in  the  manufacture  of 
fireworks.  Its  chief  use,  however,  is  in  the  manufac- 
ture of  gunpowder. 

Gunpowder. — The  value  of  gunpowder  is  due  to  the 
iact  that  it  explodes  readily,  the  explosion  being  a  chemi- 
cal change  accompanied  by  a  sudden  evolution  of  gases. 
It  is  a  mixture  of  saltpeter,  charcoal,  and  sulphur. 
When  heated,  the  saltpeter  gives  off  oxygen  and  nitro- 
gen ;  the  oxygen  combines  with  the  charcoal,  forming 
carbon  dioxide  and  carbon  monoxide ;  and  the  sulphur 
combines  with  the  potassium,  forming  potassium  sul- 
phide. When  a  mixture  of  saltpeter  and  charcoal  is 
burned,  the  reaction  which  takes  place  is  this : 

2KNO3  +  30  =  C02  +  CO  +  N2  +  K2CO3. 

By  adding  the  necessary  quantity  of  sulphur  the  car- 
bon dioxide,  which  would  otherwise  remain  in  combina- 
tion with  the  potassium  as  potassium  carbonate,  is  giveR 
off,  and  potassium  sulphide  formed  : 

2KNO3  +  30  +  S  =  3C02  +  Na  +  K2S. 


494  INORGANIC  CHEMISTRY. 

For  this  reaction  the  constituents  should  be  mixed  in 
the  proportions : 

Saltpeter, 74.83 

Charcoal,     .     .     .     . 13.31 

Sulphur, 11.86 

100.00 

This  is  approximately  the  composition  of  all  powder. 
When  gunpowder  explodes,  the  gases  formed  occupy 
about  280  times  the  volume  occupied  by  the  powder 
itself. 

Potassium  Nitrite,  KNO2,  is  formed  simply  by  heating 
the  nitrate  to  a  sufficiently  high  temperature.  The  re- 
duction is,  however,  much  facilitated  by  adding  to  the 
nitrate  some  easily  oxidized  metal,  as  lead  or  iron. 
When  the  gases  formed  by  the  action  of  arsenic  trioxide 
on  nitric  acid  are  passed  into  potassium  hydroxide,  both 
the  nitrite  and  nitrate  are  formed,  and  they  can  be  sep- 
arated by  crystallization. 

Potassium  Chlorate,  KC1O3. — The  character  of  the  reac- 
tion by  which  potassium  chlorate  is  formed  when  chlo- 
rine acts  upon  a  solution  of  potassium  hydroxide  has 
already  been  discussed  (see  p.  114).  In  the  manufac- 
ture of  the  chlorate  it  is  found  advantageous  to  make 
calcium  chlorate,  and  then  to  treat  this  with  potassium 
chloride,  when,  at  the  proper  concentration,  potassium 
chlorate  crystallizes  out,  on  account  of  the  fact  that  it  is 
less  soluble  than  the  salts  which  are  brought  together. 
The  process  in  brief  consists  in  passing  chlorine  into  a 
solution  of  calcium  hydroxide  in  which  an  excess  of  hy- 
droxide is  held  in  suspension.  The  first  action  consists 
in  the  formation  of  calcium  hypochlorite.  When  the 
solution  of  this  salt  is  boiled  it  is  decomposed,  yielding 
the  chlorate  and  chloride  : 

3Ca(OCl)2  =  Ca(O3Cl)2  +  2CaCl2. 

On  now  treating  the  solution  with  potassium  chloride 
the  following  reaction  takes  place  : 


Ca(O3Cl)3  +  2KC1  =  2KC103  +  CaCl 


POTASSIUM  PERCHLOBATE.  495 

Potassium  chlorate  crystallizes  in  lustrous  crystals  of 
tlie  monoclinic  system.  Its  taste  is  somewhat  like  that 
of  saltpeter.  It  melts  at  a  comparatively  low  tempera- 
ture (334°),  and  at  352°  begins  to  decompose,  with  evolu- 
tion of  oxygen.  At  ordinary  temperatures  100  parts  of 
water  dissolve  6  parts  of  the  salt,  and  at  the  boiling  tem- 
perature 60  parts.  In  consequence  of  the  ease  with 
which  it  gives  up  its  oxygen,  the  chlorate  is  an  excellent 
oxidizing  agent,  and  it  is  constantly  used  in  this  capacity 
in  the  laboratory.  Its  oxidizing  action  is  well  illus- 
trated by  grinding  a  very  little  of  it  in  a  mortar  with  a 
little  sulphur,*  when  an  explosion  takes  place.  With 
phosphorus  the  action  is  exceedingly  violent. 

The  chief  uses  of  potassium  chlorate  are  for  the  prep- 
aration of  oxygen,  and  in  the  manufacture  of  matches 
and  fireworks.  The  tips  of  Swedish  safety  matches  are 
made  of  potassium  chlorate  and  antimony  sulphide. 
The  surface  upon  which  they  are  rubbed  to  ignite  them 
contains  red  phosphorus.  The  chlorate  is  frequently 
used  in  medicine,  particularly  as  a  gargle  in  cases  of 
sore  throat. 

Potassium  Perchlorate,  KC1O4,  is  formed  in  the  first 
stage  of  the  decomposition  of  the  chlorate  by  heat,  as  was 
explained  under  Oxygen  (which  see).  It  is  prepared 
best  by  heating  the  chlorate  in  an  open  vessel  until,  af- 
ter having  been  liquid,  it  begins  to  get  solid  again.  As 
the  salt  is  difficultly  soluble  in  water,  the  residue  is  pow- 
dered and  washed  with  water  to  remove  the  chloride, 
and  then  crystallized  from  water.  Owing  to  the  difficult 
solubility  of  this  salt  it  is  utilized  in  chemical  analysis 
for  detecting  the  presence  of  potassium.  For  this  pur- 
pose a  solution  of  perchloric  acid  is  added  to  the  solu- 
tion under  examination,  and,  if  a  precipitate  is  formed, 
the  presence  of  potassium  may  be  inferred.  When 
heated  to  about  400°  the  perchlorate  gives  up  its  oxygen, 
and  is  reduced  to  the  chloride.  It  is  used  to  some  ex- 
tent in  the  manufacture  of  fireworks,  instead  of  the  chlo- 


*  Great  care  should  be  taken  with  all  experiments  with  potassium 
chlorate.     See  description  of  experiments. 


496  INORGANIC  CHEMISTRY. 

rate,  which,  owing  to  its   greater   instability,  is    more 
dangerous. 

Potassium  Periodate,  KIO4,  is  formed  by  the  action  of 
chlorine  on  a  mixture  of  potassium  hydroxide  and  po- 
tassium iodate.  As  has  been  stated  in  discussing 
the  acids  of  iodine,  this  salt  is  only  one  form  of  a 
group  of  potassium  salts  called  periodates,  all  of 
which  are  closely  related  to  the  normal  acid,  I(OH)7. 
Among  these  salts,  for  example,  are  the  mesoperiodate, 
K3IO6  +  4H2O,  and  the  diperiodate,  K4I2O9  +  9H2O.  The 
former  is  a  salt  of  the  acid  H3IO6,  which  is  derived  from 
the  normal  acid  by  loss  of  water,  thus  : 


Diperiodic  acid  is  derived  from  the  normal  acid  as 
represented  in  this  equation  : 

2I(OH),  =  H.I.O.  +  6H.O. 

Potassium  Cyanide,  KCN.  —  Under  Cyanogen  it  was 
stated  that  when  nitrogen  is  passed  over  a  highly  heated 
mixture  of  carbon  and  potassium  carbonate,  potassium 
cyanide  is  formed  ;  and  that  carbon  containing  nitrogen 
compounds,  as  animal  charcoal,  when  ignited  with  potas- 
'sium  carbonate,  reduces  the  carbonate,  forming  potas- 
sium, in  presence  of  which  the  carbon  and  nitrogen 
combine,  forming  the  cyanide.  The  simplest  way  to 
make  the  cyanide  is  by  heating  potassium  ferrocyanide, 
K4Fe(CN)6,  which  is  the  starting-point  in  the  prepara- 
tion of  all  cyanogen  compounds.  It  breaks  down  first 
into  potassium  cyanide  and  ferrous  cyanide,  thus  : 

K4Fe(CN)6  =  4KCN  +  Fe(CN)2. 

The  ferrous  cyanide  is,  however,  decomposed  by  heat 
into  free  nitrogen  and  carbide  of  iron, 


so  that  the  complete  decomposition  of  the  ferrocyanide 
is  represented  by  this  equation  : 

K4Fe(CN)6  =  4KCN  +  FeC2  +  N3. 


POTASSIUM  CYANIDE.  497 

As  part  of  the  cyanogen  is  lost  in  this  operation,  potas- 
sium carbonate  is  commonly  added  to  the  ferrocyanide. 
This  acts  upon  the  ferrous  cyanide,  forming  potassium 
cyanide  and  ferrous  carbonate  : 

Ee(CN)2  +  K2C03  =  FeCO3  +  2KCN. 

But  the  ferrous  carbonate  breaks  down  under  the  influ- 
ence of  heat  into  ferrous  oxide  and  carbon  dioxide 

FeCO3  =  FeO  +  CO2 ; 

and  the  ferrous  oxide  then  gives  up  its  oxygen  to  a  part 
of  the  potassium  cyanide,  converting  it  into  the  cyanate. 
The  complete  reaction  between  the  ferrocyanide  and  car- 
bonate is  therefore  represented  as  follows : 

K4Fe(CN)6  +  K2C03  =  5KCN  +  KCNO  +  CO2  +  Fe. 

The  product  obtained  in  this  way  necessarily  contains 
some  of  the  cyanate,  but  for  ordinary  purposes  this  does 
no  harm.  Potassium  cyanide  is  extremely  easily  soluble 
in  water,  and  is  deliquescent  in  moist  air.  When  boiled 
with  water  it  is  decomposed,  forming  potassium  formate 
and  ammonia : 

KCN  +  2H2O  =  KC02H  +  NH3. 

It  combines  readily  with  oxygen  when  in  the  molten  con- 
dition, as  shown  in  its  action  upon  ferrous  oxide  and 
upon  lead  oxide  (see  p.  404).  In  consequence  of  this 
power  to  combine  with  oxygen  to  form  the  cyanate,  it  is 
a  valuable  reducing  agent,  and  is  not  unfrequently  used 
in  the  laboratory  in  this  capacity. 

Just  as  the  fluoride,  chloride,  bromide,  and  iodide  of 
potassium  combine  with  the  fluorides,  chlorides,  bro- 
mides, and  iodides  of  the  metallic  elements  in  general,  so 
potassium  cyanide  combines  with  the  cyanides  of  the 
metallic  elements,  forming  the  double  cyanides.  These 
have  generally  a  composition  analogous  to  that  of  the 
double  chlorides  and  of  similar  compounds.  Thus,  silver 
cyanide  forms  the  compound  AgCy.KCy  or  AgKCy2,  in 
which  the  cyanogen  group  CN  is  represented  by  the 
symbol  Cy,  as  is  customary ;  ferrous  cyanide  forms  the 


498  INORGANIC  CHEMISTRY. 

compounds  K4FeCy6,  or  4KCy.FeCy2,  and  K3FeCy6,  or 
3KCy.FeCy3.  These  double  cyanides  are  for  the  most 
part  soluble  in  water  ;  hence  potassium  cyanide  dissolves 
many  deposits  of  metallic  salts.  It  is  frequently  used 
in  the  laboratory  in  analytical  operations. 

Potassium  Cyanate,  KCNO. — The  cyanate  is  formed  by 
oxidation  of  the  cyanide,  when  this  is  melted  and  lead 
oxide  or  minium  added  to  the  molten  mass.  It  is  most 
easily  prepared  by  heating  together  potassium  ferrocya- 
nide  and  manganese  dioxide.  The  action  consists  first 
in  the  decomposition  of  the  ferrocyanide,  and  the  subse- 
quent oxidation  of  the  potassium  cyanide  thus  formed. 
The  mass  is  extracted  with  alcohol,  as  the  cyanate  is  de- 
composed by  water  even  at  the  ordinary  temperature, 
the  products  being  potassium  carbonate  and  ammonium 
carbonate  : 

2KCNO  +  4H,O  =  K2C03  +  (NH4)2CO3. 

Acids  set  cyanic  acid  free  from  the  cyanate,  but  the 
acid  is  at  once  decomposed  by  water,  thus : 

CNOH  +  H2O  =  CO2  -f  NH3. 

Potassium  Sulphocyanate,  KCNS. — Just  as  potassium 
cyanide  takes  up  oxygen  to  form  the  cyanate,  it  also 
takes  up  sulphur  to  form  the  sulphocyanate  : 

KCN  +  S  =  KCNS. 

It  is  easily  prepared  by  adding  sulphur  to  molten 
potassium  cyanide,  or  by  heating  a  mixture  of  dehy- 
drated potassium  ferrocyanide,  potassium  carbonate, 
and  sulphur.  It  crystallizes  particularly  well  out  of  its 
solution  in  alcohol.  It  is  deliquescent,  and  when  dis- 
solved in  water  it  causes  a  very  considerable  lowering 
of  the  temperature.  Thus,  when  500  grams  of  the  salt 
are  mixed  with  400  grams  of  water  at  the  ordinary  tem- 
perature, the  temperature  sinks  to  about  —  20°.  Unlike 
the  cyanate,  it  is  not  decomposed  by  water. 

Potassium  Sulphate,  K2SO4. — This  salt  occurs  in  com- 
bination with  others  in  nature,  particularly  in  the  mineral 


PRIMARY,  OR  ACID,  POTASSIUM  SULPHATE.        499 

kainite,  which  contains  the  constituents  of  potassium 
sulphate,  magnesium  sulphate,  and  magnesium  chloride, 
as  indicated  in  the  formula  K2SO4.MgSO4.MgCl2  +  6H2O. 
This  occurs  in  Stassfurt  and  in  Kalusz.  Potassium  sul- 
phate is  used  in  medicine,  and  in  the  preparation  of 
ordinary  alum  and  of  potassium  carbonate. 

Primary,  or  Acid,  Potassium  Sulphate,  KHSO4. — This 
salt  is  obtained  as  a  secondary  product  in  the  prepara- 
tion of  nitric  acid  by  the  action  of  sulphuric  acid  upon 
saltpeter.  It  occurs  in  nature  in  the  Grotto  del  Solfo, 
near  Naples ;  and  is  made  by  treating  the  neutral  salt 
with  concentrated  sulphuric  acid.  When  heated  above 
its  melting  point  it  gives  off  water,  and  is  transformed 
into  the  disulphate,  thus  : 

2KHS04  =  K2S207  +  H20. 

When  the  disulphate  is  heated  in  contact  with  basic 
oxides  it  breaks  down,  forming  sulphates.  The  decom- 
position is  that  represented  in  the  equation 

K2S207  =  K2S04  +  SO,. 

The  nascent  sulphur  trioxide  thus  set  free  acts  with  great 
energy  upon  the  oxides  which  are  present.  Hence  acid 
potassium  sulphate  is  a  valuable  agent  for  the  purpose 
of  decomposing  some  mineral  substances  which  do  not 
readily  yield  to  the  ordinary  reagents.  Its  action  consists 
in  breaking  down  into  the  disulphate  and  water,  the  disul- 
phate then  further  breaking  down  into  normal  sulphate 
and  sulphur  trioxide.  Besides  the  salt  just  mentioned, 
which  is  known  as  the  disulphate  or  pyrosulphate,  there 
are  some  other  salts  known,  which  are  derived  from  an- 
other form  of  disulphuric  acid.  Two  good  examples 
are  the  salts  represented  by  the  formulas  K3H(SO4)2  and 
KH3(SO4)2.  The  acid  from  which  these  are  derived  has 
the  formula  H4S2O8.  It  is  to  be  regarded  as  derived 
from  two  molecules  of  normal  sulphuric  acid  by  elimina- 
tion of  four  molecules  of  water  : 


500  INORGANIC  CHEMISTRY. 

The  acid  probably  has  the  constitution  represented  thus  : 

(HO)sOS<g>SO(OH),  =  H&0r 

Sulphites.  —  When  sulphur  dioxide  is  passed  into  a 
solution  of  potassium  carbonate  until  carbon  dioxide 
ceases  to  escape,  potassium  sulphite,  K2SO3,  is  formed.  If 
the  gas  is  passed  to  saturation  the  product  is  the  pri- 
mary or  acid  sulphite,  KHSO3.  If  the  solution  of  the 
carbonate  is  hot  and  concentrated,  the  product  is  the 
disulphite,  K2S2O6,  which  bears  to  the  sulphite  the  same 
relation  that  potassium  disulphate  bears  to  the  sulphate. 
It  is  the  salt  of  an  acid  of  the  formula  H2S2O5,  which  is 
disulphurous  acid  : 


This  bears  to  sulphurous  acid  the  same  relation  that  di- 
sulphuric  acid  bears  to  sulphuric  acid. 

Carbonates.  —  The  normal  salt,  K2CO3,  is  the  chief  con- 
stituent of  wood-ashes.  When  these  are  extracted  with 
water  the  carbonate  passes  into  solution  and  the  salt 
thus  obtained  can  be  purified  in  a  number  of  ways.  The 
impure  salt  is  known  as  potash.  Formerly  all  the  potas- 
sium carbonate  made  was  obtained  from  wood-ashes,  but 
at  present  not  more  than  half  of  the  supply  comes  from 
this  source.  The  other  sources  are  the  residues  from 
the  manufacture  of  beet-sugar,  potassium  sulphate  and 
chloride,  and  wool-fat.  The  preparation  of  the  carbon- 
ate from  the  sulphate  and  chloride  is  accomplished  by 
the  same  method  as  that  used  in  the  preparation  of 
sodium  carbonate  from  the  chloride.  The  methods  used 
for  this  purpose  will  be  treated  of  under  the  head  of 
Sodium  Carbonate  (which  see).  The  salt  crystallizes 
from  very  concentrated  solutions  in  water.  It  is  deli- 
quescent, and  dissolves  in  water  with  an  evolution  of 
heat,  and  the  solution  has  a  strong  alkaline  reaction. 

Acid  Potassium  Carbonate,  HKCO3,  is  formed  by  pass- 
ing carbon  dioxide  over  the  normal  salt,  or  into  the  con- 
centrated aqueous  solution  of  the  latter.  It  is  much  less 
easily  soluble  in  water  than  the  normal  salt.  The  dry 


RUBIDIUM— CESIUM.  501 

salt  gives  off  carbon  dioxide  and  water  easily  when 
heatedj  and  is  converted  into  the  normal  salt : 

2KHC03  =  K2CO3  +  CO2  +  H2O. 

The  same  decomposition  takes  place  when  the  water 
solution  is  heated,  and  even  on  evaporation  at  the  ordi- 
nary temperature. 

Phosphates. — Three  phosphates  of  potassium  are 
known :  (1)  Tertiary,  or  normal  potassium  phosphate, 
K3PO4 ;  (2)  secondary,  or  di-potassium  phosphate,  K2HPO4 ; 
and  (3)  primary,  or  mono-potassium  phosphate,  KH2PO4. 
There  is  nothing  particularly  characteristic  about  these 
salts,  except  the  decompositions  which  the  primary  and 
secondary  salts  undergo  when  heated.  These  decompo- 
sitions have  already  been  referred  to  (see  p.  329  and  p. 
480). 

Potassium  Silicate,  K2SiO3. — A  compound  of  the  definite 
composition  represented  by  the  formula  here  given  has 
not  been  prepared.  A  solution  of  potassium  silicate  in 
water  is  prepared  by  dissolving  sand  or  amorphous  sili- 
con dioxide  in  potassium  carbonate  or  hydroxide.  It  k 
prepared  on  the  large  scale  by  melting  together  quartz 
powder  and  purified  potash.  It  is  known  as  water  glass, 
for  the  reason  that  its  solution  dries  in  the  air,  forming 
a  glass-like  looking  mass.  To  distinguish  it  from  the 
water  glass  made  with  sodium  carbonate  or  hydroxide 
it  is  called  potash  water  glass. 

BUBIDIUM,  Eb  (At.  Wt.  84.78). 
OESIUM,  Cs  (At.  Wt.  131.89). 

Both  these  elements  are  widely  distributed,  but  only 
in  small  quantities.  They  generally  occur  in  company 
with  potassium,  which  they  resemble  closely.  They 
were  discovered  by  means  of  the  spectroscope  by  Bun- 
sen  and  Kirchhoff.  The  characteristic  spectrum  of 
rubidium  consists  of  two  dark  red  lines,  and  this  is  the 
origin  of  the  name  rubidium  (from  rubidus,  dark  red). 
Csesium  was  found  in  the  Durkheim  mineral  water,  and 
was  recognized  by  two  characteristic  blue  lines,  and  the 
name  caesium  was  given  to  it  on  this  account  (from  ccesius, 


502  INORGANIC  CHEMISTRY. 

sky-blue).  Rubidium  is  found  in  different  varieties  of 
mica,  known  as  lepidolite.  The  mineral  pollux,  which  is 
essentially  a  silicate  of  caesium  and  aluminium,  contains 
caesium  as  one  of  the  chief  constituents. 

It  is  a  remarkable  fact  that  the  elements  rubidium 
and  caesium  which  are  so  similar  to  potassium  accom- 
pany it  so  generally  in  nature.  Similar  facts  were  noted 
in  the  group  consisting  of  chlorine,  bromine,  and  iodine, 
and  that  of  sulphur,  selenium,  and  tellurium.  It  will 
be  remembered  that  chlorine  is  frequently  accompanied 
by  bromine  and  iodine ;  and  sulphur  by  selenium  and 
tellurium ;  but  that  chlorine  and  sulphur  are  present  in 
much  larger  quantities  than  the  elements  which  accom- 
pany them.  Further,  the  relations  between  the  atomic 
weights  of  the  members  of  each  group  are  approximately 
the  same. 

Rubidium  is  prepared  by  the  same  method  as  that 
used  in  the  preparation  of  potassium. 

'It  is  silver-white  with  a  yellowish  tint.  It  can  be  con- 
verted into  vapor  which  has  a  blue  color.  It  takes  fire 
in  the  air  at  the  ordinary  temperature.  Its  action  upon 
water  is  the  same  as  that  of  potassium,  and  its  salts  are 
very  similar  to  those  of  potassium. 

Caesium  has  not  yet  been  isolated.  By  subjecting  the 
chloride  to  the  action  of  a  powerful  electric  current 
globules  of  metal  are  given  off  at  one  of  the  poles,  but 
these  take  fire  in  contact  with  the  air  at  the  ordinary 
temperature.  The  salts  of  caesium  are  much  like  those 
of  rubidium  and  potassium. 

SODIUM,  Na  (At.  Wt.  22.82). 

Occurrence. — Sodium  occurs  very  widely  distributed 
and  in  large  quantities  in  nature,  principally  as  sodium 
chloride.  It  is  found  in  a  number  of  silicates,  and  is  a 
constituent  of  plants,  especially  of  those  which  grow  in 
the  neighborhood  of  the  sea-shore  and  in  the  sea.  Just 
as  the  ashes  of  inland  plants  are  rich  in  potassium  car- 
bonate, so  the  ashes  of  sea  plants  and  those  which  grow 
near  the  sea  are  rich  in  sodium  carbonate.  It  is  found 
evervwhere  in  the  soil,  but  generally  in  small  quantities. 


PREPARATION  OF  SODIUM.  503 

Its  presence  in  the  soil  is  due  to  the  decomposition  of 
minerals  containing  it,  such  as  soda  feldspar,  or  albite.  It 
occurs  also  as  sodium  nitrate  or  Chili  saltpeter,  and  in 
large  quantity  in  Greenland  in  the  form  of  cryolite,  which, 
as  has  been  explained,  is  a  so-called  double  fluoride  of 
aluminium  and  sodium,  of  the  formula  Na3AlF6,  or 
AlF3.3NaF. 

Preparation. — It  is  prepared  from  sodium  carbonate  by 
the  same  method  as  that  used  in  the  preparation  of  potas- 
sium, the  reaction  involved  being  represented  thus  ; 

Na2C03  +  20  =  2Na  +  SCO. 

The  reduction  takes  place  more  readily  than  in  the  case 
of  potassium,  and  it  is  not  necessary  to  prepare  the  mix- 
ture of  carbonate  and  charcoal  by  heating  the  salt  of  an 
organic  acid,  as  is  done  in  the  preparation  of  potassium. 
The  carbonate  is  mixed  with  charcoal,  or  powdered  an- 
thracite coal,  and  calcium  carbonate,  and  sometimes  this 
mass  is  mixed  with  an  oil  and  then  ignited  in  a  crucible. 
A  successful  method  for  the  preparation  of  sodium 
on  the  large  scale  has  been  devised  by  Castner.  This 
consists  essentially  in  the  reduction  of  sodium  hydroxide 
by  heating  it  with  an  intimate  mixture  of  finely  divided 
iron  and  carbon.  The  mass  is  prepared  by  mixing  the 
iron  with  molten  pitch,  allowing  it  to  cool,  breaking  it 
into  pieces,  and  heating  to  a  comparatively  high  temper- 
ature without  access  of  air.  The  reduction  is  said  to 
take  place  at  a  temperature  of  825°,  instead  of  1400°  as 
in  the  older  method.  The  main  reaction  is  represented 
by  this  equation : 

GNaOH  +  FeC2  =  2Na2CO3  +  6H  +  2Na  +  Fe. 

The  preparation  of  sodium  was  formerly  of  more 
importance  than  it  is  at  present,  for  the  separation  of 
aluminium  from  the  compounds  found  in  nature  depended 
upon  the  preparation  of  sodium.  Methods,  depending 
upon  the  use  of  an  electric  furnace,  have  been  devised 
for  the  preparation  of  aluminium,  and  the  old  method 
involving  the  use  of  sodium  is  no  longer  employed. 


504  INORGANIC  CHEMISTRY. 

Properties. — The  properties  of  sodium  are  very  similar 
to  those  of  potassium.  It  is  light,  floating  on  water ;  it 
has  a  bright  metallic  lustre ;  and  at  the  ordinary  tem- 
perature it  is  soft  like  wax.  It  decomposes  water,  but 
not  as  readily  as  potassium  does.  Its  specific  gravity  is 
0.9735;  its  melting  point  95.6°.  Its  vapor  is  colorless 
when  seen  in  thin  layers,  while  thick  layers  appear  pur- 
ple. When  melted  and  allowed  to  cool  it  takes  the 
crystallized  form.  When  exposed  to  the  air  it  acts  upon 
the  moisture,  and  is  converted  into  the  hydroxide. 

Applications. — It  is  used  for  the  purpose  of  isolating 
some  elements  whose  oxides  cannot  easily  be  reduced, 
as,  for  example,  aluminium,  magnesium,  and  silicon, 
which  are  prepared  by  treating  their  chlorides  with 
sodium.  Silicon,  however,  as  we  have  seen,  is  prepared 
better  by  treating  potassium  fluosilicate,  K2SiF6,  with 
sodium.  The  element  is  also  used,  in  combination  with 
mercury  as  sodium  amalgam,  a  substance  which  affords 
a  ready  means  of  making  nascent  hydrogen.  It  also 
finds  constant  application  in  the  laboratory  for  a  variety 
of  purposes. 

Sodium  Hydride,  Na2H,  is  formed  in  the  same  way  as 
the  corresponding  compound  of  potassium,  and  is  in 
every  way  similar  to  it. 

Sodium  Chloride,  NaCl. — This  is  the  substance  which 
is  generally  known  simply  as  salt,  or  common  salt.  It 
occurs  very  widely  distributed,  and  in  immense  quantities, 
in  the  earth.  The  most  important  deposits  are  those  at 
Wieliczka  in  Galicia,  at  Stassfurt  and  Keichenhall  in  Ger- 
many, and  at  Cheshire  in  England.  Besides  these  there 
are,  however,  many  other  deposits  in  the  United  States 
of  America,  in  Africa,  and  in  Asia.  As  it  is  easily  soluble 
in  water,  many  springs  and  streams,  as  well  as  lakes  and 
the  ocean,  contain  it.  Sea- water  contains  2.7  per  cent. 
In  some  places  sodium  chloride  is  taken  out  of  mines  in 
solid  form.  Frequently,  however,  water  is  allowed  to 
flow  into  cavities  in  the  earth,  and  to  remain  for  some 
time  in  contact  with  the  salt.  The  solution  thus  formed 
is  afterward  drawn  or  pumped  out  of  the  mine  and 
evaporated  by  appropriate  methods.  It  is  generally 


SODIUM  CHLORIDE.  505 

allowed  slowly  to  run  down  walls  made  of  twigs,  so  that 
a  large  surface  of  the  liquid  is  exposed  to  the  air.  The 
concentrated  solution  thus  obtained  is  then  evaporated 
to  crystallization  by  the  aid  of  heat. 

In  hot  countries  salt  is  obtained  by  the  evaporation  of 
sea- water,  the  heat  of  the  sun  being  used  for  the  purpose. 
Large  shallow  cavities  are  made  in  the  earth,  and  into 
these  the  water  flows  at  high-tide,  or  it  is  pumped  up  into 
them  if  they  are  too  high.  The  process  is  continued  for 
some  months,  and  then  the  mother-liquor  is  drawn  off, 
and  the  accumulated  salt  collected  and  subjected  to 
proper  methods  of  purification. 

The  salt  obtained  by  the  above  methods  is  not  pure. 
It  always  contains  sodium  sulphate,  together  with  mag- 
nesium and  calcium  chlorides.  The  chlorides  of  magnes- 
ium and  calcium  cause  it  to  become  moist  in  the  air. 
Pure  salt  does  not  attract  moisture. 

Sodium  chloride  crystallizes  in  colorless  and  trans- 
parent cubes.  Some  of  that  which  occurs  in  nature  has 
a  blue  color.  When  deposited  from  an  evaporating 
solution  it  takes  the  form  of  small  cubes  arranged  in 
groups  of  the  shape  of  hollow  pyramids,  known  as  the 
hopper-shaped  deposits.  If  urea  or  boric  acid  is  present 
in  the  solution  the  crystals  of  sodium  chloride  are  octa- 
hedrons or  combinations  of  these  with  cubes.  When 
deposited,  the  crystals  enclose  water,  not  as  water  of 
crystallization,  and  this  is  given  off  when  the  crystals 
are  heated,  the  action  being  accompanied  by  a  crackling 
sound.  This  is  known  as  decrepitation. 

Sodium  chloride  melts  at  776°,  and  is  volatile  at  a  red 
heat.  In  hot  water  it  is  but  little  more  soluble  than  in 
cold.  At  100°  100  parts  of  water  dissolve  39  parts,  and 
at  ordinary  temperatures  36  parts. 

Sodium  chloride  is  the  starting-point  in  the  preparation 
of  all  sodium  compounds,  as  well  as  of  chlorine  and  hy- 
drochloric acid.  Salt  is  necessary  to  the  life  of  man 
and  many  other  animals.  The  role  played  by  it  in  the 
animal  economy  is  not  understood,  but  it  is  found 
generally  distributed  throughout  the  body  in  small 
quantity. 


506  INORGANIC  CHEMISTRY. 

The  fluoride,  bromide,  and  iodide  of  sodium  are  like  the 
corresponding  potassium  salts  and  need  not  be  described. 

Sodium  Hydroxide,  NaOH. — This  compound  resembles 
potassium  hydroxide  in  all  respects.  Being  cheaper  it  is 
used  much  more  extensively.  It  is  prepared  in  the  same 
way,  by  treating  sodium  carbonate  in  solution  with  cal- 
cium hydroxide,  when  insoluble  calcium  carbonate  and 
soluble  sodium  hydroxide  are  formed  : 

Na2CO3  +  Ca(OH)2  =  CaCO3  +  2NaOH. 

The  substance  is  commonly  called  caustic  soda.  It  is 
extensively  used  for  the  purpose  of  making  soap  from 
fats. 

Oxides. — Sodium  forms  two  oxides,  the  monoxide, 
Na2O,  and  the  peroxide,  Na2O2.  In  this  respect  a  differ- 
ence is  noticed  between  sodium  and  potassium  ;  the  latter 
forming  the  compounds  K2O  and  K2O4. 

Sodium  Peroxide,  Na2O2,  has  acquired  importance  in 
the  arts  as  a  bleaching-agent.  It  is  prepared  by  heating 
sodium  in  a  current  of  dry  air  at  a  temperature  of  300°. 
When  heated  to  a  high  temperature  it  gives  off  oxygen. 
Water  decomposes^ it,  forming  sodium  hydroxide,  and 
setting  oxygen  free. 

The  hydrosulphide  and  the  sulphides  of  sodium  are  made 
just  as  the  potassium  compounds  are,  and  resemble  them 
very  closely. 

Sodium  Sulphantimonate,  NasSbS4,  also  known  as 
Schlippe's  salt,  is  a  particularly  beautiful  example  of  the 
salts  of  sulpho-acids.  It  is  made,  as  its  composition  indi- 
cates, by  dissolving  antimony  pentasulphide  in  a  solution 
of  sodium  sulphide : 

Sb2S6  +  3Na2S  =  2Na3SbS4. 

Sodium  Nitrate,  NaNO3. — This  compound  occurs  in 
large  quantity  in  southern  Peru  on  the  border  of  Chili, 
and  is  known  as  Chili  saltpeter.  The  natural  salt  con- 
tains, besides  the  nitrate,  sodium  chloride,  sulphate,  and 
iodate.  Sodium  nitrate  is  very  similar  to  potassium 
nitrate,  but  it  cannot  be  used  in  place  of  the  more  ex- 
pensive potassium  salt  in  the  manufacture  of  the  finer 


SODIUM  SULPHATE.  507 

grades  of  gunpowder,  as  it  becomes  moist  in  the  air,  and 
does  not  decompose  quickly  enough.  It  is  used  ex- 
tensively in  the  manufacture  of  nitric  acid,  and  also  for 
the  purpose  of  preparing  ordinary  saltpeter.  The  iodine 
contained  in  Chili  saltpeter  is  now  extracted  on  the  large 
scale,  and  this  forms  an  important  source  of  iodine. 

Sodium  Sulphate,  Na2SO4. — This  salt  was  first  made  by 
Glauber,  as  it  is  now  made,  by  the  action  of  sulphuric 
acid  on  sodium  chloride.  It  is  commonly  called  Glauber's 
salt.  It  occurs  in  a  number  of  natural  waters,  as  in  that 
of  Friedrichshall  and  Carlsbad.  It  occurs,  further,  in 
solid  form  in  small  quantities  in  some  localities.  It  is 
made  in  very  large  quantities  in  connection  with  the 
manufacture  of  soda,  the  first  reaction  in  this  process  con- 
sisting in  treating  sodium  chloride  with  sulphuric  acid. 
It  is  also  formed  in  the  manufacture  of  nitric  acid  by 
the  action  of  sulphuric  acid  on  Chili  saltpeter. 

Large  quantities  of  sodium  sulphate  are  now  made  by 
the  action  of  magnesium  sulphate  on  sodium  chloride. 
This  process  is  employed  at  Stassfurt,  where  both  mag- 
nesium sulphate  and  sodium  chloride  occur  in  immense 
quantities.  The  action  takes  place  between  concentrated 
solutions  at  low  temperatures.  It  is  represented  by  the 
equation 

2NaCl  +  MgS04  =  Na2SO4  +  MgCl2. 

It  crystallizes  in  large,  colorless,  monoclinic  prisms, 
which  contain  ten  molecules  of  water.  These  crystals 
are  formed,  however,  only  in  case  the  temperature  of  the 
solution  is  below  33°  at  the  time  they  are  deposited.  If 
a  saturated  solution  is  cooled  down  to  a  point  some- 
where between  33°  and  40°,  the  salt  is  deposited  without 
water  of  crystallization.  When  the  crystallized  salt  is 
heated  to  33°  it  loses  a  part  of  its  water.  The  salt  is  most 
easily  soluble  in  water  at  33°  ;  above  this  point  the  solu- 
bility decreases.  Taking  these  facts  into  consideration, 
it  appears  probable  that  in  solutions  below  33°  the  com- 
pound Na2SO4  -f-  10H2O  is  present ;  while  if  the  solu- 
tion is  heated  above  this  point  the  compound  breaks 


508  INORGANIC  CHEMISTRY. 

down,  and  the  anhydrous  salt,  as  well  as  the  salts  with 
less  than  ten  molecules  of  water,  are  less  easily  soluble. 
One  of  the  ten  molecules  of  water  is  held  in  the  com- 
pound more  firmly  than  the  rest.  It  seems  probable 
that  this  is  not  present  as  water  but  as  hydroxyl,  the  salt 

having  the  formula  OS  j  ^\   (=Na2SO4  +  H2O). 

Sodium  sulphate  easily  forms  supersaturated  solutions 
which  crystallize  rapidly  if  disturbed,  if  a  small  crystal 
of  the  salt  is  thrown  into  them,  and  if  cooled  down 
to  —  8°.  This  phenomenon  is  frequently  presented  by 
salts,  but  it  is  shown  in  a  particularly  striking  way  by 
this  one. 

When  expose^d  to  the  air  the  salt  loses  its  water  of 
crystallization  and  crumbles  to  a  white  powder.  This  is 
the  process  already  described  as  efflorescence  (see  p.  58). 

Sodium  sulphate  is  used  as  a  purgative  in  medicine, 
and  in  the  laboratory  for  the  production  of  cold  arti- 
ficially. A  good  freezing  mixture  is  made  by  bringing  it 
together  with  concentrated  hydrochloric  acid.  Sodium 
chloride  is  formed,  and  the  water  of  crystallization  of  the 
sulphate  takes  the  liquid  form.  This  change  from  the 
solid  to  the  liquid  form  is  accompanied  by  a  marked  ab- 
sorption of  heat.  Ice  can  be  made  in  this  way  without 
difficulty.  The  chief  uses  of  the  sulphate  are  in  'the 
manufacture  of  sodium  carbonate  and  of  glass,  as  will 
be  explained  farther  on. 

Sodium  Thiosulphate,  Na2S2O3  +  5H2O. — This  is  the  salt 
which  is  commonly  called  hyposulphite  of  soda.  It  is 
made  on  the  large  scale  by  treating  caustic  soda  with 
sulphur,  and  conducting  sulphur  dioxide  into  the  solu- 
tion. As  has  been  pointed  out,  when  sulphur  acts  upon 
potassium  carbonate  polysulphides  of  potassium  and  the 
thiosulphate  are  formed.  A  similar  action  takes  place 
when  sulphur  acts  upon  caustic  soda.  The  polysul- 
phides in  the  solution  give  up  sulphur  to  the  sulphite 
and  convert  it  thus  into  the  thiosulphate  : 

Na,S,  +  Na2SO3  =  Na2S  +  Na2S8O3. 


SODIUM  CARBONATE.  509 

It  is  also  made  by  boiling  a  solution  of  sodium  sulphite 
and  adding  sulphur : 

NaaSO3  +  S  =  Na2S2O3. 

Its  chief  application  is  in  photography,  in  which  art  it  is 
used  for  the  purpose  of  dissolving  the  excess  of  silver 
salt  on  the  plate  which  has  been  exposed  to  the  light, 
and  on  which  a  picture  has  been  developed.  The  action 
consists  in  the  formation  of  salts  in  which  both  sodium 
and  silver  are  contained.  These  are  soluble  in  water. 
The  thiosulphate  will  be  taken  up  more  in  detail  under 
Silver  (which  see). 

Sodium  Carbonate,  Na2CO3. — This  salt,  commonly  called 
soda,  is  one  of  the  most  important  of  manufactured  chemi- 
cal compounds.  The  mere  mention  of  the  fact  that  it  is 
essential  to  the  manufacture  of  glass  and  soap  will  serve 
to  give  some  conception  of  its  importance.  It  is  found 
in  the  ashes  of  sea  plants,  just  as  potassium  carbonate  is 
found  in  the  ashes  of  inland  plants.  Formerly,  it  was 
made  entirely  from  plant  ashes,  but  we  are  no  longer  de- 
pendent upon  this  source  for  our  supply  of  the  salt,  as 
two  methods  have  been  devised  for  preparing  it  from 
sodium  chloride,  with  which  nature  provides  us  in  such 
abundance.  As  these  methods  are  of  great  importance, 
and  are,  further,  very  interesting  applications  of  chemi- 
cal principles,  they  will  be  described  below. 

Properties. — Anhydrous  sodium  carbonate  is  a  powder 
which  is  formed  by  heating  the  crystallized  salt.  It  melts 
to  a  clear  liquid  when  heated  to  a  sufficiently  high  tem- 
perature. It  dissolves  in  water  very  readily  with  evolution 
of  heat.  The  action  is,  however,  not  as  marked  as  in 
the  case  of  potassium  carbonate.  When  the  salt  is 
deposited  from  a  water  solution  it  has  the  composition 
Na2CO3  +  10H2O.  This  salt,  it  will  be  observed,  con- 
tains the  same  number  of  molecules  of  water  of  crystal- 
lization as  sodium  sulphate.  Like  this,  too,  it  effloresces 
when  exposed  to  the  air.  When  heated  it  melts  in  its 
water  of  crystallization,  and  the  salt  Na2CO3  -f-  H2O,  or 
(HO)2C(ONa)2,  separates.  This,  however,  loses  water 


510  INORGANIC  CHEMISTRY. 

when  heated  higher,  and  is  converted  into  the  anhydrous 
salt.  The  conduct  of  the  carbonate  towards  water  at  dif- 
ferent temperatures  is  suggestive  of  that  of  the  sulphate. 
Its  maximum  solubility  is  at  temperatures  between  38° 
and  70°.  Above  the  latter  point  the  solubility  decreases. 
The  cause  of  this  phenomenon  is,  in  all  probability, 
the  same  as  that  referred  to  in  describing  the  analogous 
phenomenon  presented  by  the  sulphate ;  that  is,  the  ex- 
istence of  the  hydrated  compound  Na2CO3  -|-  10H2O  in 
solution  at  temperatures  below  70°,  and  the  dissociation 
of  this  compound  into  water  and  salts  containing  a  smaller 
number  of  molecules  of  water  of  crystallization,  which 
are  less  soluble,  when  the  temperature  is  raised  above 
this  point.  The  crystals  of  sodium  carbonate  containing 
ten  molecules  of  water  of  crystallization  belong  to  the 
monoclinic  system. 

Applications. — Sodium  carbonate  is  used  in  immense 
quantities  in  the  manufacture  of  glass,  and  in  the  prepa- 
ration of  caustic  soda,  which  is  used  in  the  manufacture 
of  soap. 

The  lie  Blanc  Process  for  the  Manufacture  of  Sodium 
Carbonate. — In  the  manufacture  of  soda  the  problem  to 
be  solved  is  to  convert  sodium  chloride  into  sodium  car- 
bonate. The  first  method  devised  for  this  purpose  is 
that  of  Le  Blanc.  During  the  French  revolution  the 
supply  of  potash  was  cut  off  from  France.  This  led  the 
government  to  offer  a  prize  for  a  practical  method  for 
manufacturing  soda  from  common  salt.  The  method 
proposed  by  Le  Blanc  at  that  time,  and  which,  until  re- 
cently, has  been  used  almost  exclusively  involves  three 
reactions : 

(1)  The    sodium   chloride  is  converted  into  sodium 
sulphate  by  treating  it  with  sulphuric  acid : 

2NaCl  +  H2S04  =  Na2SO4  +  2HC1. 

(2)  The  sodium  sulphate  thus  obtained  is  heated  with 
charcoal,  which  reduces  it  to  sodium  sulphide  : 

Na2SO4  +  20  =  Na3S  +-2CO,. 


SODIUM  CARBONATE- LE  BLANC  PROCESS.        511 

(3)  The  sodium  sulphide  is  heated  with  calcium  car- 
bonate, when  sodium  carbonate  and  calcium  sulphide 
are  formed : 

Na2S  +  CaCO3  =  NaaCO3  +  CaS. 

The  conversion  of  the  sulphate  into  the  carbonate  is, 
therefore,  expressed  by  the  equation 

Na2SO4  +  20  +  CaCO3  =  Na2CO3  +  CaS  +  2CO2. 

Calcium  sulphide  is  insoluble  in  water,  so  that  by 
treating  the  resulting  mass  with  water  the  sodium  car- 
bonate is  separated  from  the  sulphide. 

In  practice  the  sodium  sulphate  is  mixed  with  coal 
and  calcium  carbonate,  and  the  mixture  heated  in  ap- 
propriately constructed  furnaces.  The  coal  reduces  the 
sulphate  to  sulphide,  which  then  reacts  upon  the  cal- 
cium carbonate  as  above  represented.  The  product  of 
the  action  is  known  as  crude  soda  or  black  ash.  It  con- 
tains, as  its  chief  constituents,  sodium  carbonate  and 
calcium  sulphide,  together  with  some  calcium  oxide,  and 
a  number  of  other  substances  in  small  quantities.  In 
order  to  purify  this  product,  it  is  broken  to  pieces,  and 
treated  with  water ;  and  the  solution  thus  obtained 
evaporated,  when  the  salt  of  the  composition  Na.2CO3  -(- 
2H2O  is  deposited.  This  is  dipped  out,  and  dried  by 
heat,  when  it  loses  all  its  water.  The  product  is  the 
calcined  purified  soda  of  commerce.  This  always  contains 
some  sulphate  and  chloride  together  with  a  small  quan- 
tity of  sulphite. 

When  dissolved  in  water  and  allowed  to  crystallize, 
the  salt  is  deposited  in  large  crystals  which  contain 
water  in  the  proportion  represented  by  the  formula 
Na2CO3  -f-  10H2O.  This  is  the  so-called  crystallized  soda. 

Most  of  the  soda  which  comes  into  the  market  is  the 
calcined  variety.  The  mother-liquors  from  the  crystal- 
lized soda  contain  some  sodium  hydroxide  in  consequence 
of  the  action  of  calcium  hydroxide  on  sodium  carbonate. 
This  can  be  converted  into  soda  by  passing  carbon  di- 


512  INORGANIC  CHEMISTRY. 

oxide  into  it ;  and  it  can  also  be  partly  separated  from 
the  carbonate  and  brought  into  the  market  as  such. 

A  method  has  recently  been  devised  for  the  purpose  of 
avoiding  the  manufacture  and  use  of  sulphuric  acid  in 
the  soda  factories.  This  consists  in  passing  a  hot  mixture 
of  sulphur  dioxide,  air,  and  steam  over  sodium  chloride. 
The  action  which  takes  place  is  represented  by  this 
equation : 

2NaCl  +  SOa  +  H2O  +  O  =  NaaS04  +  2HC1. 

As,  in  the  manufacture  of  soda,  by  the  Le  Blanc  pro- 
cess, the  sulphur  remains  in  combination  as  calcium 
sulphide,  a  process,  known  as  the  Chance  process,  has 
been  devised  for  its  recovery.  This  consists  in  passing 
carbon  dioxide  into  the  waste,  thus  liberating  hydrogen 
sulphide ;  passing  this  into  another  portion  of  the  waste, 
thus  converting  the  calcium  sulphide  into  the  hydro- 
sulphide  ;  and  then  treating  this  with  carbon  dioxide, 
when  a  gas  rich  in  hydrogen  sulphide  is  given  off : 

COa  +  CaHaSa  +  HaO  =  CaCO3  +  2HaS. 

By  regulating  the  supply  of  air  the  gas  is  burned  either 
to  sulphur  dioxide  or  to  sulphur. 

Ammonia  Process  for  the  Manufacture  of  Soda. — An- 
other process  now  in  extensive  use  for  the  manufacture 
of  soda  is  the  so-called  ammonia  process,  or  the  Solvay 
process.  This  depends  upon  the  fact  that  mono-sodium 
carbonate,  HNaCO3,  is  comparatively  difficultly  soluble 
in  water.  If,  therefore,  mono-ammonium  carbonate,  or 
acid  ammonium  carbonate,  HNH4CO3,  is  added  to  a  solu- 
tion of  common  salt,  acid  sodium  carbonate,  HNaCO3, 
crystallizes  out,  and  ammonium  chloride  remains  in  the 
solution : 

NaCl  +  HNH4CO3  =  HNaC03  +  NH4C1. 
When  the  acid  carbonate  thus  obtained  is  heated,  it  gives 


SODA  FROM  CRYOLITE.  513 

off  carbon  dioxide,  and  is  converted  into  the  normal  salt 
thus  : 

2HNaCO3  =  Na2CO3  +  CO2  +  H2O. 

The  carbon  dioxide  given  off  is  passed  into  ammonia, 
and  thus  again  obtained  in  the  form  of  acid  ammonium 
carbonate : 

NH3  +  H2O  +  CO2  =  HNH4CO3. 

The  ammonium  chloride  obtained  in  the  first  reaction 
is  treated  with  lime  or  magnesia,  MgO,  and  the  ammonia 
set  free.  This  ammonia  is  used  again  in  the  preparation 
of  acid  ammonium  carbonate.  The  object  of  using  mag- 
nesia is  to  get  magnesium  chloride,  which,  when  evap- 
orated to  dryness  and  heated,  yields  magnesia  and  hy- 
drochloric acid : 

MgCl2  +  H20  =  MgO  +  2HC1. 

More  than  half  the  soda  supply  of  the  world  is  now  fur- 
nished by  the  Solvay  process. 

Manufacture  of  Soda  from  Cryolite. — As  cryolite  occurs 
in  nature  in  large  quantities,  and  can  be  obtained  cheaply, 
it  is  used  in  some  places  for  the  manufacture  of  soda. 
The  reactions  involved  are  : 

(1)  The  action  of  calcium  carbonate  upon  cryolite  at 
a  high  temperature,  when   sodium  aluminate,   calcium 
fluoride,  and  carbon  dioxide  are   formed  as  represented 
in  the  equation 

Na3AlF6  +  3CaCO3  =  3CaF2  +  Na3AlO3  +  3CO2. 

(2)  The  action  of  carbon  dioxide  upon  the  solution  of 
the  aluminate,  when  aluminium  hydroxide  is  precipitated, 
and  sodium  carbonate  formed  which  remains  in  solution  : 

2Na3A103  +  3C02  +  3H2O  =  3Na2CO3  +  2A1(OH)3. 

After  the  mixture  of  cryolite  and  calcium  carbonate,  or 
chalk,  has  been  heated,  the  mass  is  treated  with  water, 
when  the  sodium  aluminate  dissolves,  while  the  calcium 
fluoride  does  not.  After  separating  the  solution  from 
the  insoluble  residue,  carbon  dioxide  is  passed  through  it. 


514  INORGANIC  CHEMISTRY. 

Mono-Sodium  Carbonate,  Primary  Sodium  Carbonate, 
HNaCOg. — This  salt  is  commonly  called  "  bi-carbonate  of 
soda."  It  is  easily  prepared  by  passing  carbon  dioxide 
over  the  ordinary  carbonate  dissolved  in  its  water  of 
crystallization : 

NaaC03  +  COa  +  H2O  =  2HNaCO3. 

When  heated  it  gives  up  carbon  dioxide  and  water,  and 
is  converted  into  the  normal  salt.  As  was  stated  in  con- 
nection with  the  ammonia-soda  process,  primary  sodium 
carbonate  is  much  more  difficultly  soluble  in  water  than 
the  normal  salt.  At  ordinary  temperatures  100  parts  of 
water  dissolve  about  10  parts  of  the  salt. 

It  is  used  in  medicine,  and  extensively  in  the  prepara- 
tion of  soda-water  and  other  effervescing  drinks. 

Sodium-Potassium  Carbonate,  KNaCO3  +  12H2O,  is  an 
interesting  example  of  a  salt  of  a  dibasic  acid  containing 
two  different  metals.  It  is  easily  made  by  mixing  solu- 
tions of  potassium  and  sodium  carbonates,  and  is  ob- 
tained in  the  form  of  large  crystals. 

Phosphates. — There  are  three  phosphates  of  sodium 
just  as  there  are  three  phosphates  of  potassium.  The 
point  of  chief  interest  presented  by  them  is  that  the 
secondary  salt,  HNa2PO4,  is  the  one  most  easily  obtained, 
and  is  the  substance  commonly  known  as  sodium  phos- 
phate. When  a  solution  of  this  salt  is  treated  with  an 
excess  of  sodium  hydroxide,  and  the  solution  evaporated, 
normal  or  tertiary  sodium  phosphate  crystallizes  out.  This 
has  the  composition  Na3PO4  +  12H2O.  The  solution  of 
the  latter  salt  has  an  alkaline  reaction,  and  when  ex- 
posed to  the  air  it  absorbs  carbon  dioxide,  and  is  con- 
verted into  the  secondary  salt : 

2Na3P04  +  CO,  +  H2O  =  2HNa2PO4  +  Na2C03. 

Secondary  sodium  phosphate,  HNa2PO4  -f-  12H2O,  is 
easily  made  by  adding  sodium  carbonate  to  a  solution  of 
phosphoric  acid  until  an  alkaline  reaction  is  shown.  It 
is  also  prepared  on  the  large  scale  from  bone-ash.  It 
forms  monoclinic  prisms  which  effloresce  in  the  air. 


SODIUM  BORATE.  515 

Sodium  Metaphosphate,  NaPO3,  is  formed  when  the  pri- 
mary phosphate  is  ignited.  There  are  several  modifica- 
tions of  the  salt  which  appear  to  differ  from  one  another 
as  represented  in  the  formulas  NaPO3,  Na2P2O6,  Na3P3O9, 
etc.  This  relation  is  called  polymerism;  or  substances 
which  have  the  same  composition  but  different  molecu- 
lar weights  are  said  to  be  polymeric.  Relations  of  this 
kind  are  very  common  among  the  compounds  of  carbon. 
Among  the  hydrocarbons  mentioned  in  Chapter  XIX, 
for  example,  are  acetylene,  C2H2,  and  benzene,  C6H,. 
There  are,  further,  two  other  hydrocarbons  of  the  for- 
mulas C4H4  and  C8H8.  Plainly  these  hydrocarbons  all 
have  the  same  percentage  composition.  They  are  poly- 
meric in  the  sense  in  which  that  term  has  been  defined. 

Di-sodium  Pyro-antimonate,  H2Na2Sb2O7  +  6H2O,  is  of 
special  interest  because  it  is  insoluble  in  cold  water,  and 
may  therefore  be  used  for  the  purpose  of  detecting  so- 
dium in  analysis.  It  is  formed  when  a  solution  of  the 
corresponding  potassium  salt  is  added  to  a  solution  of 
a  sodium  salt. 

Sodium  Borate. — Normal  boric  acid,  as  we  have  seen, 
has  the  composition  B(OH)3,  and  there  are  a  number  of 
borates  derived  from  this  acid  by  direct  replacement  of 
the  hydrogen  by  metals.  The  salt  which  boric  acid 
most  readily  forms  with  sodium  hydroxide  or  sodium 
carbonate,  however,  is  that  derived  from  tetraboric  acid, 
H2B4O7,  which  is  derived  from  normal  boric  acid  by 
elimination  of  water.  (See  p.  355).  This  salt  is  borax, 
which  in  crystallized  form  has  the  composition  repre- 
sented by  the  formula  Na2B4O7  +  10H2O.  By  adding 
the  required  quantity  of  sodium  hydroxide  to  a  solution 
of  borax,  and  evaporating  to  crystallization,  sodium 
metaborate,  NaBO2  +  4H2O,  is  obtained  : 

Na2B407  +  2NaOH  =  4NaBO2  +  HaO. 

The  metaborate  is  decomposed  when  its  solution  is  ex- 
posed to  the  action  of  the  air.  It  is  thus  converted  by 
carbon  dioxide  into  sodium  carbonate  and  borax,  or 
sodium  tetraborate. 


516  INORGANIC  CHEMISTRY. 

Borax  occurs  in  nature  in  several  lakes  in  Asia  and 
in  Clear  Lake,  Nevada,  in  the  United  States.  It  is  man- 
ufactured by  neutralizing,  with  sodium  carbonate,  the 
boric  acid  found  in  Tuscany.  When  heated,  borax  puffs 
up,  and  at  red  heat  it  melts,  forming  a  transparent,  color- 
less liquid.  The  dehydrated  salt  is  known,  as  anhydrous 
or  calcined  borax.  In  the  molten  condition,  borax  has 
the  power  to  combine  with  metallic  oxides,  and,  as  many 
of  the  double  borates  thus  formed  are  colored,  the  salt 
is  used  in  blow-pipe  work  for  the  purpose  of  detecting 
certain  metals.  As  it  dissolves  metallic  oxides,  it  is 
used  in  the  process  of  soldering,  as  it  is  necessary 
to  have  bright,  untarnished  metallic  surfaces  in  order 
that  the  solder  shall  adhere  firmly.  The  action  of  mol- 
ten borax  upon  metallic  oxides  is  similar  to  that  which 
takes  place  when  sodium  hydroxide  acts  upon  a  solution 
of  borax.  Borates  of  the  metals  are  formed  together 
with  sodium  borate,  or  double  borates  in  which  part  of 
the  hydrogen  is  replaced  by  sodium  and  part  by  other 
metals. 

Borax  is  extensively  used  in  the  manufacture  of  por- 
celain and  in  glass-painting.  It  is  an  antiseptic,  pre- 
venting the  decomposition  of  some  organic  substances. 

Sodium  Silicate,  Na2SiO3. — Sodium  silicate  is  formed 
by  dissolving  silicon  dioxide  in  sodium  hydroxide,  and 
can  be  obtained  in  crystallized  form.  It  is  prepared  on 
the  large  scale  by  melting  together  quartz  sand  and  so- 
dium carbonate  in  the  proper  proportions,  andfby  melt- 
ing together  sodium  sulphate,  quartz  sand,  and  charcoal 
powder.  This  substance  is  commonly  known  as  water- 
glass.  It  is  soluble  in  water,  and,  when  its  solution 
dries,  it  leaves  a  transparent  coating  on  the  surface  on 
which  it  is  placed.  It  is  extensively  used  in  the  manu- 
facture of  artificial  stone. 

LITHIUM,  Li  (At.  Wt.  .6.97). 

Lithium  occurs  in  nature  in  relatively  small  quantity, 
chiefly  in  the  minerals  lepidolite,  petalite,  and  spodu- 
mene,  and  in  many  mineral  waters.  It  is  also  found  in 


AMMONIUM  SALTS.  517 

the  ashes  of  a  number  of  plants.  It  is  prepared  by  the 
electrolysis  of  the  chloride  in  the  molten  condition. 
The  metal  is  silver- white,  and  is  characterized  by  its  low 
specific  gravity.  It  acts  vigorously  upon  water,  but,  if 
the  water  is  at  the  ordinary  temperature,  the  hydrogen 
given  off  does  not  take  fire.  In  the  air  it  conducts  itself 
in  much  the  same  way  that  sodium  does. 

The  most  characteristic  salts  of  lithium  are  the  phos- 
phate, carbonate,  and  chloride. 

Lithium  Phosphate,  Li3PO4  +  pI2O,  is  precipitated  when 
secondary  sodium  phosphate  is  added  to  a  solution  of  a 
lithium  salt.  It  is  very  difficultly  soluble  in  water  at  the 
ordinary  temperature. 

Lithium  Carbonate,  Li2CO3,  is  also  rather  difficultly  sol- 
uble in  water,  and  is  deposited  when  a  solution  of  sodi- 
um carbonate  is  added  to  a  fairly  concentrated  solution 
of  lithium  chloride.  It  dissolves  uric  acid,  which  is  in- 
soluble in  water,  and  is  therefore  used  in  medicine  for 
the  purpose  of  removing  pathological  deposits  of  this 
acid  in  the  body.  For  this  purpose  it  is  generally  ad- 
ministered in  the  form  of  a  solution  in  water  containing 
carbon  dioxide. 

Lithium  Chloride,  LiCl,  is  peculiar  on  account  of  the 
fact  that  it  is  soluble  in  alcohol  and  in  a  mixture  of  al- 
cohol and  ether.  In  this  respect  it  differs  from  the 
chlorides  of  potassium  and  sodium,  which  are  insoluble 
in  alcohol.  If,  therefore,  a  mixture  of  the  chlorides  of 
the  three  metals  is  treated  with  alcohol,  only  lithium 
chloride  dissolves ;  and  in  this  way  lithium  can  be  sep- 
arated from  the  other  metals. 

AMMONIUM  SALTS. 

Attention  has  already  been  called  to  the  marked  simi- 
larity of  the  salts  of  potassium  and  sodium  to  those 
formed  by  the  action  of  ammonia  on  the  acids,  and 
known  as  ammonium  salts.  The  most  important  of 
these  salts  will  be  briefly  considered  in  this  connection. 
A  characteristic  property  of  ammonium  salts  which  dis- 
tinguishes them  from  the  salts  of  all  the  metals  is  their 


518  INORGANIC  CHEMISTRY. 

volatility.  "When  sublimed,  they  all  undergo  decomposi- 
tion, which  is  either  partial  or  complete. 

The  simplest  kind  of  decomposition  which  they  un- 
dergo is  dissociation  into  ammonia  and  the  acid.  This 
is  illustrated  in  the  case  of  ammonium  chloride,  which, 
when  heated  to  a  sufficiently  high  temperature,  is  dis- 
sociated into  ammonia  and  hydrochloric  acid.  This  is 
an  example  of  true  dissociation.  The  amount  of  decom- 
position is  constant  for  any  given  temperature  and  pres- 
sure. 

An  ammonium  salt  of  a  polybasic  acid  containing  some 
metal  gives  off  ammonia  and  leaves  an  acid  salt,  which 
generally  undergoes  further  decomposition.  Thus,  so- 
dium-ammonium sulphate,  NaNH4SO4,  first  gives  off 
ammonia  and  forms  mono-sodium  sulphate  : 


The   acid   salt   thus  formed  then  undergoes  further 
change  and  the  pyrosulphate  is  formed  : 


Another  example  of  this  kind  of  decomposition  of  am- 
monium salts  is  that  afforded  by  sodium  -  ammonium 
phosphate,  HNaNH4PO4.  When  heated,  this  gives  off 
ammonia  and  then  water,  the  final  product  being  sodium 
metaphosphate  : 


(ONa 
PO-^  ONH4 

(OH 

(ONa 
=  POl  OH   +NH3; 

(OH 

(ONa 
PO^  OH 
OH 

=  PO2ONa     +  H2O. 

Some  ammonium  salts  undergo  deeper-seated  decom- 
positions, and  do  not  give  ammonia  as  one  of  the  prod- 
ucts. This  is  true  especially  of  such  salts  as  readily 
give  off  oxygen.  In  such  cases  the  ammonia  is  oxidized, 


AMMONIUM  SALTS.  519 

so  that  the  hydrogen  forms  water.  This  is  illustrated 
in  the  decomposition  of  ammonium  nitrate  and  nitrite  L 

NH4N03  =  N20  +  2H20  ;  and 
NH4N02:=N2    +2H20. 

Further,  all  ammonium  salts  are  decomposed  with  evo- 
lution of  ammonia  when  treated  with  basic  hydroxides. 
This  has  been  illustrated  in  the  preparation  of  ammonia 
from  ammonium  chloride  by  treatment  with  calcium 
hydroxide : 

2NH4C1  +  Ca(OH)2  =  CaCl2  +  2NH3  +  2H2O. 

The  ammonium  salts  are  made  by  neutralizing  acids 
with  ammonia. 

Ammonium  Chloride,  NH4C1. — This  salt  is  commonly 
called  sal  ammoniac.  At  present  its  principal  source  is 
the  so-called  ammoniacal  liquor  of  the  gas-works.  This 
liquid  contains  a  considerable  quantity  of  ammonium 
carbonate,  and,  when  it  is  treated  with  lime,  ammonia  is 
given  off.  This  is  passed  into  hydrochloric  acid,  and  the 
solution  of  ammonium  chloride  thus  formed  evaporated 
to  crystallization.  The  salt  has  a  sharp,  salty  taste,  and 
is  easily  soluble  in  water.  When  heated,  it  is  converted 
into  vapor  without  melting  and  with  very  slight  decom- 
position ;  and  when  the  vapor  comes  in  contact  with  a 
cold  surface,  it  condenses  in  crystalline  form.  This  pro- 
cess of  vaporizing  and  condensing  a  solid  is  called  sub- 
limation. Some  of  the  ammonium  chloride  met  with  in 
the  market  has  been  sublimed.  The  salt  is  used  in  the 
preparation  of  ammonia,  in  medicine,  and  for  other  pur- 
poses. When  it  is  dissolved  in  water,  a  considerable 
lowering  of  temperature  is  caused. 

Ammonium  Sulphocyanate,  NH4CNS. — This  salt  is  pre- 
pared by  bringing  together  aqueous  ammonia,  carbon 
disulphide,  and  alcohol.  The  first  product  is  ammonium 
thiocarbamate,  the  formation  of  which  is  perfectly  analo- 
gous to  the  formation  of  the  ordinary  carbamate  by  the 
action  of  carbon  dioxide  on  ammonia : 


520  INORGANIC  CHEMISTRY. 


The  thiocarbamate  afterwards  breaks  down  when 
heated,  forming  the  sulphocyanate  and  hydrogen  sul- 
phide : 


CS  =  CNSNH4  +  H3S. 


The  salt,  like  so  many  other  ammonium  salts,  causes 
a  marked  lowering  of  temperature  when  dissolved  in 
water.  When  100  grams  are  dissolved  in  the  same 
weight  of  water  at  17°,  the  temperature  falls  to  —12°.  It 
is  now  much  used  in  analytical  processes  for  the  esti- 
mation of  silver  and  copper. 

Ammonium  Sulphide,  (NH4)aS.  —  This  compound  is  ex- 
tensively used  in  chemical  analysis  for  the  purpose  of 
precipitating  those  sulphides  which  are  soluble  in  dilute 
hydrochloric  acid  (see  p.  198  and  p.  473).  As  will  be  re- 
membered, in  the  usual  method  of  analyzing  a  mixture 
of  substances,  the  first  step  consists  in  adding  hydro- 
chloric acid  to  the  solution.  This  precipitates  silver, 
lead,  and,  under  certain  conditions,  mercury.  This  pre- 
cipitate having  been  filtered  off,  hydrogen  sulphide  is 
passed  through  the  filtrate,  when  those  metals  are  pre- 
cipitated whose  sulphides  are  insoluble  in  dilute  hydro- 
chloric acid.  The  precipitate  is  filtered  off,  and  ammo- 
nium sulphide  added  to  the  filtrate,  when  the  metals 
whose  sulphides  are  soluble  in  dilute  hydrochloric  acid 
are  thrown  down.  Among  these  are  iron,  cobalt,  nickel, 
manganese,  etc.  Any  other  soluble  sulphide  might  be 
used  ;  but  the  advantage  of  ammonium  sulphide  is  that  it 
is  volatile,  and  hence,  by  evaporating  the  solution  and 
heating,  it  can  be  got  rid  of  after  it  has  served  its  pur- 
pose. Another  use  to  which  it  is  put  in  analysis  is  for 
the  purpose  of  dissolving  the  sulphides  of  tin,  arsenic, 
and  antimony,  which  are  precipitated  by  hydrogen  sul- 
phide, and  thus  separating  these  from  the  other  sul- 
phides of  the  group.  This  solution  depends  upon  the 


AMMONIUM  SULPHIDE.  521 

poVer  of  the  sulphides  to  form  salts  of  sulpho- acids,  as 
has  been  repeatedly  explained. 

Ammonium  sulphide  is  made  by  passing  hydrogen 
sulphide  into  an  aqueous  solution  of  ammonia.  If  the 
gas  is  passed  until  the  solution  is  saturated,  the  product 
is  the  hydrosulphide  : 

NH3  +  H2S  =  NH4HS. 

If  only  half  this  quantity  of  the  gas  is  passed,  the  pro- 
duct is  the  sulphide  : 

2NH3  +  H2S  =  (NH4)2S. 

The  simplest  way  to  make  it,  however,  is  to  divide  a 
quantity  of  a  solution  of  ammonia  into  two  equal  parts  ; 
saturate  one  half,  thus  forming  the  hydrosulphide,  and 
add  the  other  half,  when  this  reaction  takes  place  : 
HNH4S  +  NH3  =  (NH4)2S. 

The  product  is  a  colorless  liquid  of  a  disagreeable 
odor.  It  soon  changes  color,  becoming  yellow,  and  after 
a  time  a  yellow  deposit  is  formed  in  the  vessel  in  which 
it  is  contained.  This  change  of  color  is  due  to  the  action 
of  the  oxygen  of  the  air.  Some  of  the  sulphide  is  de- 
composed into  ammonia,  water,  and  sulphur,  thus  : 

(NH4)2S  +  O  =  2NH3  +  H20  +  S. 

The  sulphur  set  free  in  this  way  combines  with  the 
undecomposed  ammonium  sulphide,  forming  the  com- 
pounds  (NH4)2S2,  (NH4)2S3,  (NH4)2S4,  and  (NH4)2S5.  When 
as  much  sulphur  has  been  set  free  as  is  required  to  form 
the  pentasulphide,  further  decomposition  by  the  oxygen 
of  the  air  causes  a  deposit  of  sulphur.  Therefore,,  in 
bottles  containing  ammonium  sulphide  which  are  al- 
lowed to  stand  for  a  long  time  a  deposit  of  sulphur  is 
always  found.  A  solution  containing  the  polysulphides 
is  called  yellow  ammonium  sulphide.  It  is  this  which  is 
used  for  the  purpose  of  dissolving  the  sulphides  of 
arsenic,  antimony,  and  tin  in  analytical  operations. 

As  stated  above,  a  solution  of  ammonium  hydrosul- 
phide, HNH4S,  is  made  by  passing  hydrogen  sulphide 
into  a  solution  of  ammonia  until  no  more  is  taken  up. 


522  INORGANIC  CHEMISTRY. 

Ammonium  Nitrate,  NH4NO3,  is  obtained  in  crystals, 
which  are  easily  soluble  in  water.  It  is  of  use  chiefly  in 
the  preparation  of  nitrous  oxide.  When  heated  sud- 
denly to  a  high  temperature  it  is  decomposed  rapidly 
into  nitrogen,  water,  and  nitric  oxide  : 

2NH4N03  =  N2  +  2NO  +  4H2O. 

This  decomposition  may  take  place  in  the  preparation  of 
nitrous  oxide  if  in  the  last  stages  of  the  operation  the 
heat  is  raised  too  high,  and  explosions  may  be  caused 
in  this  way.  When  dissolved  in  water  a  marked  lower- 
ing of  temperature  takes  place. 

Ammonium  Carbonate,  (NH4)2CO3  —  When  dry  ammonia 
gas  and  dry  carbon  dioxide  are  brought  together,  they 
unite  and  form  the  salt  known  as  ammonium  carbamate, 

which  has  the  composition  CO 


This  is  the  salt  of  an  acid,  CO  j  QTTSJ  known  as  car- 

bamic  acid.     When  the  carbamate  is  dissolved  in  water, 
it  is  converted  into  the  carbonate  : 


ONH 


When  heated  to  58°,  the  normal  carbonate  is  decomposed, 
forming  carbon  dioxide,  water,  and  ammonia.  The  sub- 
stance found  in  the  market  under  the  name  of  ammoni- 
um carbonate  is  made  by  heating  together  ammonium 
chloride  or  sulphate  and  chalk.  It  consists  of  normal 
ammonium  carbonate,  (NH4)2CO3,  primary  ammonium 
carbonate,  HNH4(CO3),  and  ammonium  carbamate. 

Primary  Ammonium  Carbonate,  HNH4CO3,  is  formed  by 
treating  the  normal  carbonate  with  carbon  dioxide,  and 
by  allowing  the  commercial  carbonate  to  lie  exposed  to 
the  air,  when  the  carbamate  is  converted  into  the  car- 
bonate by  the  moisture,  and  the  carbonate  loses  am- 
monia : 


REACTIONS  OF  THE  MEMBERS  OF  THE  SODIUM  GRO  UP.  523 

It  is  easily  decomposed  into  ammonia,  water,  and  car- 
bon dioxide. 

Sodium-ammonium  Phosphate,  HNaNH4PO4. — This  salt 
is  known  as  microcosmic  salt,  and  is  much  used  in  the 
laboratory  in  blow-pipe  work.  It  is  contained  in  guano 
and  in  decomposed  urine.  It  is  easily  made  by  mixing 
solutions  of  di-sodium  phosphate  and  ammonium  chlo- 
ride, and  allowing  to  crystallize.  In  crystallized  form  it 
contains  four  molecules  of  water,  HNaNH4PO4  -f-  4H2O. 
The  changes  which  the  anhydrous  salt  undergoes  when 
heated  were  described  on  page  518.  When  the  crystal- 
lized salt  is  heated,  the  water  of  crystallization  is  first 
given  off.  The  value  of  the  salt  in  blow-pipe  work  de- 
pends upon  the  fact  that  at  high  temperatures  the  meta- 
phosphate  combines  with  metallic  oxides,  forming  mixed 
phosphates,  the  reactions  being  like  those  which  meta- 
phosphoric  acid  undergoes  with  water  : 

2HP03    +  H20  =  H4P207; 
HP03      +  H20  =  H3P04; 
2NaP03  +  M20  =  Na2M2P207 ; 
NaP03    +  M2O  =  NaM2P04. 

Many  of  these  double  phosphates  and  pyrophosphates 
are  colored,  and,  like  the  double  borates  (see  p.  516) 
they  furnish  a  means  of  detecting  some  of  the  metals. 

Reactions  of  the  Members  of  the  Sodium  Group  which 
are  of  Value  in  Chemical  Analysis. — The  chief  difficulty 
experienced  in  chemical  analysis  is  in  distinguishing 
between  similar  elements.  Sodium  and  potassium,  for 
example,  conduct  themselves  so  much  alike  in  so 
many  respects  that  we  might  subject  them  to  the  in- 
fluence of  a  number  of  reagents  without  being  able 
to  tell  which  one  we  are  working  with.  For  pur- 
poses of  analysis,  therefore,  it  is  necessary  to  take 
advantage  of  differences  between  the  elements,  and  the 
more  striking  the  differences  the  better.  Those  reac- 
tions which  give  rise  to  the  formation  of  insoluble  com- 
pounds or  precipitates  are  most  frequently  used  in 
analysis.  Very  few  salts  of  the  members  of  the  sodium 


524  INORGANIC  CHEMISTRY. 

group  are  insoluble,  and  the  difficulty  of  distinguishing 
between  these  elements  is  increased  by  this  fact.  In 
ordinary  analyses  the  elements  of  this  group  which  are 
of  most  importance  are  potassium  and  sodium,  the 
other  elements  of  the  group  being  but  rarely  met 
with.  Ammonium  compounds  are  easily  distinguished 
from  those  of  potassium  and  sodium  by  the  fact  that, 
when  treated  with  caustic  soda  or  potash,  they  give  off 
ammonia,  which  is  recognized  by  its  characteristic  odor. 
The  chief  reactions  which  are  of  value  in  distinguishing 
between  potassium  and  sodium  are  the  following : 

Platinum  Chloride,  PtCl4,  forms  difficultly  soluble  salts 
with  potassium  and  ammonium  chlorides.  These  are 
the  cUoroplatinates,  K2PtCl6  and  (NH4)9PtCle.  The  cor- 
responding salt  of  sodium  is  easily  soluble. 

Perchloric  Acid,  HC1O4,  forms  difficultly  soluble  potas- 
sium perchlorate,  KC1O4,  when  added  to  solutions  of  po- 
tassium salts. 

Fluosilicic  Acid,  H2SiF6,  forms  difficultly  soluble  salts 
with  potassium  and  sodium,  K2SiF6  and  Na2SiF6,  but  not 
with  ammonium. 

Tartaric  Acid,  H2(C4H4O6),  forms  a  difficultly  soluble 
potassium  salt  of  the  formula  KH(C4H4O6).  The  corre- 
sponding salt  of  sodium  is  easily  soluble.  The  forma- 
tion of  mono-potassium  tartrate  takes  place  as  repre- 
sented in  the  equation  : 

KOI  +  Ha(C,H,0.)  =  KH(C,H40.)  +  HC1. 

Normal  or  neutral  potassium  tartrate  is  soluble  in  water, 
so  that,  if  the  difficultly  soluble  acid  tartrate  is  filtered 
off,  and  potassium  carbonate  added  to  it,  it  dissolves  in 
consequence  of  the  formation  of  the  neutral  salt,  which 
takes  place  as  represented  in  the  equation 

2KH(C.H40.)  +  K,CO,  =  2K,(C.H,O.)  +  CO,  +  H3O. 

If,  to  the  solution  of  the  neutral  salt,  hydrochloric  acid 
is  added,  the  acid  salt  is  again  formed  and  precipitated  : 

K2(C4H4O6)  +  HC1  =  KH(C4H4O6)  +  KC1. 

Di-sodium  Pyro-antimonate,  Na2H2Sb2O7,  is  insoluble  in 
cold  water,  and  is  formed  when  a  solution  of  the  corre- 


FLAME  REACTIONS  AND  THE  SPECTROSCOPE.      525 

spending  potassium  salt  is  added  to  a  solution  of  a  so- 
dium salt. 

Flame  Reactions  and  the  Spectroscope. — When  a  clean 
piece  of  platinum  wire  is  held  for  some  time  in  the  flame 
of  the  Bunsen  burner,  it  then  imparts  no  color  to  the 
flame.  If  now  a  small  piece  of  sodium  carbonate  or  any 
other  salt  of  sodium  is  put  on  it,  the  flame  is  colored 
intensely  yellow.  All  sodium  compounds  have  this 
power,  and  hence  the  chemist  makes  use  of  this  fact  for 
the  purpose  of  detecting  the  presence  of  sodium.  Simi- 
larly, potassium  compounds  color  the  flame  violet ;  lith- 
ium compounds  color  the  flame  red ;  rubidium  and 
caesium  produce  colors  similar  to  that  of  the  potassium 
flame.  While  it  is  an  easy  matter  to  recognize  potas- 
sium alone,  or  any  one  of  the  other  metals  alone,  it  is 
difficult  to  do  so  when  they  are  together  in  the  same 
compound.  For  example,  when  sodium  and  potassium 
.are  together,  the  intense  yellow  caused  by  the  sodium 
completely  masks  the  more  delicate  violet  caused  by 
the  potassium,  so  that  the  latter  cannot  be  seen  by  the 
unaided  eye.  In  this  particular  case  the  difficulty  can 
be  got  over  by  letting  the  light  from  the  flame  pass 
through  a  blue  glass,  or  through  a  thin  vessel  of  glass 
containing  a  solution  of  indigo.  The  yellow  light  is 
thus  cut  off,  while  the  violet  light  passes  through  and 
can  be  recognized.  A  more  general  method  for  de- 
tecting the  constituents  of  light  is  by  means  of  a  prism 
of  glass.  Lights  of  different  colors,  which  are  pro- 
duced by  ether  waves  of  different  lengths,  are  turned 
out  of  their  course  to  different  extents  when  passed 
through  a  prism,  as  is  seen  when  white  sunlight  is 
passed  through  a  prism.  A  narrow  beam  of  white  light 
passing  in  emerges  as  a  band  of  various  colors,  called  its 
spectrum.  We  thus  see  that  white  light  is  made  up  of 
lights  of  different  colors  ;  or,  to  speak  in  the  language  of 
physics,  that  motion  of  the  light-ether  which  produces 
upon  the  eye  the  sensation  of  white  light  is  made  up  of  a 
number  of  motions,  each  of  which  alone  produces  upon  the 
eye  the  sensation  of  a  color.  Similarly,  we  can  determine 
what  any  light  is  composed  of.  Every  light  has  its  char- 


526  INORGANIC  CHEMISTRY. 

acteristic  spectrum.  The  light  given  off  from  any  solid 
heated  to  a  white  heat  gives  a  continuous  spectrum,  like 
that  of  the  sunlight.  An  incandescent  gaseous  substance, 
on  the  other  hand,  gives  a  spectrum  made  up  of  separate 
bands  of  color,  or  a  banded  spectrum.  The  light  produced 
by  burning  sodium,  or  by  introducing  a  sodium  com- 
pound in  a  colorless  flame,  gives  a  spectrum  consisting  of 
a  narrow  yellow  band.  The  spectrum  of  the  potassium 
flame  consists  essentially  of  two  bands,  one  red  and  one 
violet.  Further,  these  bands  always  occupy  definite  posi- 
tions relatively  to  one  another,  so  that,  in  looking  through 
a  prism  at  the  light  caused  by  potassium  and  sodium,  the 
yellow  band  of  sodium  is  seen  in  its  position,  and  the 
two  potassium  bands  in  their  proper  positions.  There 
is  therefore  no  difficulty  in  detecting  these  elements 
when  present  in  the  same  substance  or  in  the  presence 
of  other  elements  which  give  characteristic  spectra. 

The  instrument  used  for  the  purpose  of  observing  the 
spectra  of  different  lights  is  called  the  spectroscope*  It 
consists  essentially  of  a  prism  and  two  telescopes. 
Through  one  of  the  telescopes  the  light  to  be  examined 
is  allowed  to  pass  so  as  to  strike  the  prism  properly. 
The  light  emerges  from  the  other  side  of  the  prism,  and 
is  observed  through  the  other  telescope,  which  is  pro- 
vided with  lenses  for  the  purpose  of  magnifying  the 
spectrum.  By  means  of  a  third  telescope,  an  image  of  a 
scale  is  thrown  upon  the  face  of  the  prism  from  which 
the  spectrum  emerges,  and  is  reflected  thence  into  the 
observing-tube,  together  with  the  spectrum,  so  that  the 
position  of  the  bands  can  be  accurately  determined.  By 
means  of  the  spectroscope,  it  is  possible  to  detect  the 
minutest  quantities  of  some  elements,  and,  since  it  was 
devised,  several  new  elements  have  been  discovered 
through  its  aid  ;  as,  for  example,  csesium,  rubidium,  thal- 
lium, indium,  gallium,  and  others. 

*  For  an  account  of  the  spectroscope  and  its  uses,  the  student  should 
consult  some  work  on  physics.  The  principles  involved  in  its  construc- 
tion and  application  are  physical  principles,  and  cannot  properly  be 
taken  up  in  detail  in  a  text-book  of  chemistry. 


CHAPTER  XXVI. 

ELEMENTS  OF  FAMILY  II,   GROUP  A: 
GLUCINUM  —  MAGNESIUM— CALCIUM— STRONTIUM- 
BARIUM  [ERBIUM]. 

General. — The  elements  of  this  group  fall  into  two  sub- 
groups. Calcium,  strontium,  and  barium  are  strikingly 
alike.  They  also  have  some  points  in  common  with  the 
members  of  the  potassium  family,  and  at  the  same  time 
are  related  in  some  degree  to  the  metals  of  Family  III, 
Group  A,  which  are  known  as  the  earth  metals.  There- 
fore, calcium,  barium,  and  strontium  are  generally  called 
the  metals  of  the  alkaline  earths.  Glucinum  and  mag- 
nesium resemble  the  metals  of  the  alkaline  earths  in  some 
ways,  but  they  also  resemble  the  members  of  Group  B, 
of  the  same  family,  which  includes  zinc  and  cadmium. 
On  comparing  the  group  with  the  elements  presented  in 
the  last  chapter,  some  analogous  facts  are  noticed.  Ar- 
ranging the  five  elements  of  the  potassium  group  in  the 
order  of  their  atomic  weights,  and  the  elements  of  Family 
II,  Group  A,  in  the  same  way,  we  have  this  table  : 


Li 

Na 

K 

Eb 

Cs 

6.97 

22.82 

38.82 

84.78 

131.89 

Gl 
9.01 

Mg 
24.10 

Ca 
39.76 

Sr 
86.95 

Ba 
136.39 

As  regards  the  analogies  between  the  elements  in  each 
group,  the  general  statement  can  be  made  that  the  last 
three  members  of  each  group  resemble  one  another  more 
closely  than  they  resemble  the  first  two  members  of  the 
group,  while  the  first  two  members  in  each  group  also 
resemble  each  other  closely.  The  natural  grouping 
according  to  the  properties  is  into  the  sub-groups  : 

(527) 


528  INORGANIC  CHEMISTRY. 


a  b 

Lithium,  Potassium, 

Sodium,  and  Rubidium, 

Caesium. 

GluciBum,  Calcium, 

Magnesium,       and  Strontium, 

Barium. 

The  relations  between  the  atomic  weights  of  the  ele- 
ments of  Family  II,  Group  A,  are  similar  to  those  of  the 
elements  of  Family  I,  Group  A.  That  of  magnesium, 
24.10,  is  nearly  half  the  sum  of  those  of  glucinum,  9.01, 
and  calcium,  39.76.  We  have 

9.01  +  39.  T6 


2 


=  24.38. 


So,  also,  that  of  strontium,  86.95,  is  approximately  half 
the  sum  of  those  of  calcium,  39.76,  and  barium,  136.39  • 

39.76  +  136.39 


In  the  calcium  group  the  specific  gravities  increase  in 
the  order  of  the  atomic  weights  : 

At.  Wt.  Sp.  Gr. 
Calcium,.     .     .     .     39.76  1.57 

Strontium,    .     .     .     86.95  2.5 

Barium,  ....  136.39  3.75 

All  the  elements  of  the  group  are  bivalent.     The  general 
formulas  of  the  principal  compounds  are  as  follows  : 

MC12,  M(OH)2,  M(N03)2,  MS04,  M3(PO4)2,  MSiO3,  etc. 

The  chlorides,  hydroxides,  and  nitrates  are  soluble  in 
water.  The  sulphates  decrease  in  solubility  as  the 
atomic  weights  increase.  Glucinum  sulphate,  G1SO4  ,  is 
soluble  in  its  own  weight  of  water  ;'  magnesium  sulphate, 
MgSO4,  is  soluble  in  about  three  times  its  weight  of 
water  ;  calcium  sulphate,  CaSO4,  dissolves  in  400  parts  ; 
strontium  sulphate,  SrSO4,  in  about  8000  parts;  and 


CALCIUM:— OCCURRENCE— PREPARATION.          529 

barium  sulphate,  BaSO4 ,  in  about  400,000  parts  of  water. 
Barium  sulphate,  as  will  be  seen,  is  practically  insoluble 
in  water.  The  normal  carbonates  of  all  except  glucinum 
are  insoluble  in  water.  The  solubility  of  the  hydroxides 
increases  as  the  atomic  weight  increases.  Glucinum  hy- 
droxide is  insoluble ;  magnesium  hydroxide  is  but 
slightly  soluble.  One  hundred  parts  of  water  at  the 
ordinary  temperature  dissolve  0.1368  parts  of  calcium  hy- 
droxide, 2  parts  of  strontium  hydroxide,  and  3.5  parts 
of  barium  hydroxide.  The  solubility  of  strontium  and 
barium  hydroxides  is,  however,  much  increased  at  higher 
temperatures. 

CALCIUM  SUB-GROUP. 

This  sub-group,  as  has  been  stated,  consists  of  the 
three  very  similar  elements,  calcium,  strontium,  and 
barium.  Of  these  calcium  occurs  most  abundantly  in 
nature.  Barium  and  strontium  frequently  accompany 
each  other,  and  both  are  found  in  some  localities  in  com- 
pany with  calcium.  They  are  much  less  abundant  in 
nature  than  calcium. 

CALCIUM,  Ca  (At.  Wt.  39.76X 

Occurrence. — Calcium  is  found  in  nature  in  enormous 
quantities,  chiefly  in  the  form  of  the  carbonate,  CaCO3, 
as  limestone,  marble,  and  chalk.  It  also  occurs  in  the  form 
of  the  sulphate,  CaSO4,  as  gypsum ;  of  the  phosphate, 
Ca3(PO4)2,  as  phosphorite  and  apatite ;  of  the  fluoride, 
CaF2,  as  fluor-spar.  It  is  found  in  solution  in  most 
natural  waters  either  as  the  carbonate  or  sulphate ;  and 
in  the  organs  of  plants  and  animals.  Bones  contain  a 
large  proportion  of  calcium  phosphate ;  egg-shells  and 
coral  contain  calcium  carbonate. 

Preparation. — The  element  is  made  by  decomposing 
molten  calcium  chloride  by  means  of  the  electric  current ; 
and  by  first  making  zinc-calcium  and  distilling  off  the 
zinc  by  heating  to  a  high  temperature  in  a  crucible  made 
of  carbon  from  a  gas-retort.  The  zinc-calcium  is  made 
by  melting  together  a  mixture  of  calcium  chloride,  zinc, 


530  INORGANIC  CHEMISTRY. 

and  sodium.  The  sodium  decomposes  the  chloride,  and 
the  reduced  metal  dissolves  in  or  combines  with  the  zinc 
as  soon  as  it  is  formed. 

Properties. — It  is  a  brass-yellow,  lustrous  metal,  which 
in  moist  air  becomes  covered  with  a  layer  of  hydroxide 
and  carbonate.  At  ordinary  temperatures  it  decomposes 
water  just  as  potassium  and  sodium  do,  but  heat  is  not 
evolved  rapidly  enough  to  set  fire  to  the  hydrogen. 
Heated  to  a  high  temperature,  it  burns  in  the  air,  forming 
the  oxide.  It  is  not  made  in  quantity,  and  has  found  no 
practical  application. 

Calcium  Chloride,  CaCL. — This  salt  is  found  in  nature 
in  combination  with  other  chlorides,  particularly  in  the 
mineral  tachydrite,  which  occurs  in  the  salt  deposits  at 
Stassfurt,  and  has  the  composition  represented  by  the 
formula  CaCl2.MgCl2  -(-  12H2O.  It  is  also  found  in  solu- 
tion in  sea- water.  It  is  obtained  as  a  by-product  in  the 
preparation  of  ammonia  from  ammonium  chloride  and 
lime ;  in  the  preparation  of  potassium  chlorate  from  cal- 
cium chlorate  and  potassium  chloride  (see  p.  494);  and  in 
the  ammonia-soda  process.  It  is  made  by  dissolving 
calcium  carbonate  in  hydrochloric  acid,  as  in  the  prepa- 
ration of  carbon  dioxide.  From  very  concentrated 
solutions  it  crystallizes  with  six  molecules  of  water, 
CaCl2  -J-  6H2O.  ,When  these  crystals  are  exposed  to  the 
air  they  soon  deliquesce.  When  a  solution  of  calcium 
chloride  is  evaporated,  and  care  is  taken  to  keep  the 
temperature  below  200°,  it  solidifies,  forming  a  porous 
mass  which  has  the  composition  represented  by  the  for- 
mula CaCl2  -f-  2H2O.  This  is  much  used  in  laboratories 
as  a  drying  agent,  as  it  absorbs  water  with  great  ease.  If 
this  salt  is  heated  above  200°  it  loses  all  its  water,  and 
the  dehydrated  chloride  melts,  forming  fused  calcium 
chloride.  This  is  also  much  used  on  account  of  its  dry- 
ing power.  Gases  are  passed  through  tubes  filled  with 
granulated  calcium  chloride  for  the  purpose  of  drying 
them,  and  the  salt  is  also  placed  in  vessels  in  which  it  is 
necessary  that  the  air  should  be  dry,  as  in  balance-cases, 
desiccators,  etc.  The  fused  salt  generally  has  a  slight 
alkaline  reaction,  which  is  caused  by  the  presence  of  a 


COMPOUNDS  OF  CALCIUM.  531 

small  quantity  of  lime.  This  is  formed  by  the  action  of 
steam  at  high  temperature  on  the  chloride,  the  reaction 
being  represented  by  this  equation  : 

CaCl2  +  H2O  =  CaO  +  2HC1. 

This  decomposition  takes  place  only  to  a  slight  extent. 
The  porous  chloride,  which  contains  two  molecules  of 
water,  does  not  contain  any  hydroxide,  and  it  is  therefore 
better  adapted  for  use'  in  cases  in  which  it  is  necessary 
that  it  should  not  absorb  carbon  dioxide,  as  in  the  analysis 
of  organic  compounds. 

Calcium  chloride  forms  crystallized  compounds  with 
ammonia  and  with  alcohol,  as  well  as  with  water.  It  is 
obvious  from  this  that  calcium  chloride  cannot  be  used 
for  the  purpose  of  drying  ammonia  gas.  When  the  com- 
pounds with  ammonia  and  with  alcohol  are  heated  they 
break  down,  yielding  ammonia  and  aksohol  respectively, 
as  the  compound  with  water  gives  up  the  latter. 

Calcium  Fluoride,  CaF2. — This  compound  occurs  in 
large  quantities  in  nature  as  the  mineral  fluor-spar.  It 
occurs  beautifully  crystallized  in  cubes,  and  is  insoluble 
in  water.  It  is  the  source  of  fluorine  compounds  in  gen- 
eral, and  is  used  in  metallurgical  operations  for  the 
reason  that  it  melts  readily  and  does  not  act  upon  other 
substances  easily.  It  therefore  simply  serves  as  a  liq- 
uid medium  in  which  reactions  take  place  at  high  tem- 
peratures. A  substance  which  acts  in  this  way  and  is 
used  for  this  purpose  is  called  a  flux.  The  name  fluor- 
spar has  its  origin  in  this  use  of  the  substance.  A  flux 
plays  to  some  extent  the  same  part  at  elevated  tempera- 
ture in  facilitating  reactions  that  water  plays  at  ordinary 
temperatures. 

Calcium  Oxide,  CaO. — This  important  compound  is 
commonly  called  lime,  or,  to  distinguish  it  from  the  hy- 
droxide or  slaked  lime,  it  is  called  quick-lime.  It  is  made 
in  large  quantity  by  heating  calcium  carbonate  in  ap- 
propriately constructed  furnaces,  known  as  lime-kilns. 
Pure  lime  is  made  by  decomposing  some  pure  form  of 
calcium  carbonate,  as  marble  or  calc-spar.  The  decom- 
position of  calcium  carbonate  is  not  complete  in  an  at- 


532  INORGANIC  CHEMISTRY. 

mosphere  of  carbon  dioxide,  hence  precautions  must  be 
taken  to  remove  the  gas  formed  by  the  decomposition. 
Further,  when  lime  is  heated  to  a  temperature  consider- 
ably higher  than  that  necessary  to  effect  the  first  decom- 
position it  again  absorbs  carbon  dioxide. 

Lime  is  a  white,  amorphous,  infusible  substance.  When 
heated  in  the  flame  of  the  compound  blow-pipe  it  gives 
an  intense  light,  as  any  other  infusible  substance  would 
do  under  the  same  circumstances.  When  exposed  to  the 
air  it  attracts  moisture  and  carbon  dioxide,  and  is  con- 
verted into,  the  carbonate.  It  must  therefore  be  protected 
from  the  air.  Lime  which  has  been  converted  into  the 
carbonate  by  exposure  to  the  air  is  said  to  be  air-slaked. 

Calcium  Hydroxide,  Ca(OH)2.  —  When  calcium  oxide  or 
quick-lime  is  treated  with  water  it  becomes  hot  and  crum- 
bles to  a  fine  powder.  The  substance  which  is  formed 
in  this  operation  is  somewhat  soluble  in  water,  the  solu- 
tion being  known  as  lime-water.  The  chemical  change 
which  takes  place  when  lime  is  treated  with  water  has 
been  explained.  It  consists  in  the  formation  of  a  com- 
pound of  the  formula  Ca(OH)2,  known  as  slaked  lime  ; 
and  the  operation  is  known  as  slaking.  The  action  is  of 
the  same  kind  as  that  with  which  we  have  so  frequently 
had  to  deal  in  the  transformation  of  oxides  into  the  cor- 
responding hydroxides.  Thus  when  potassium  oxide  is 
treated  with  water  it  is  changed  to  the  hydroxide,  with 
a  marked  evolution  of  heat,  the  reaction  being  repre- 
sented in  this  way  : 


So,  too,  when  sulphur  trioxide  is  brought  in  contact  with 
water  it  appears  to  form  the  hydroxide,  normal  sulphuric 
acid  : 


=  S- 


OH 
OH 
OH 
OH 
OH 
OH 


CALCIUM  HYDROXIDE.  533 

The  action  in  the  case  of  calcium  oxide  is  represented  in 
a  similar  way  : 


The  hydroxide  is  a  fine  white  powder.     At  red  heat  it 
loses  water  and  is  reconverted  into  the  oxide  : 


When  lime-water  is  exposed  to  the  air  it  becomes  cov- 
ered with  a  crust  of  calcium  carbonate,  and  finally  all  the 
calcium  is  precipitated  as  calcium  carbonate.  A  solution 
of  calcium  hydroxide  affords  a  convenient  means  of  de- 
tecting the  presence  of  carbon  dioxide,  as  has  been  shown 
in  dealing  with  this  gas.  The  solution  has  an  alkaline 
reaction,  and  acts  in  many  respects  like  the  hydroxides 
of  potassium  and  sodium.  Attention  has  been  called  to 
the  fact  that  the  hydroxides  of  most  of  the  metals  are 
insoluble  in  water,  and  that  when  a  soluble  hydroxide  is 
added  to  the  salt  of  such  a  metal  the  insoluble  hydrox- 
ide is  precipitated.  The  same  kind  of  decomposition  of 
salts  is  effected  by  a  solution  of  calcium  hydroxide. 
Thus,  when  it  is  added  to  ferric  chloride,  ferric  hydrox- 
ide is  thrown  down  : 


Cl  pQ/OH  (OH 

Fe    Cl            <OH  Fe^  OH 

Fe    Cl  n  ^OH  Fe-|  OH 

Cl  Ua^OH  OH 


This  reaction  is  entirely  analogous  to  that  which  takes 
place  between  ferric  chloride  and  potassium  hydroxide : 

(Cl       KOH  (OH       KC1 

Fe^  Cl  +  KOH  =  Fe  \  OH  +  KC1 . 
Cl       KOH  OH       KC1 


534  INORGANIC  CHEMISTRY. 

Lime  is  extensively  used  in  the  arts,  generally  in  the 
form  of  the  hydroxide.  As  we  have  seen,  it  is  used  in 
the  preparation  of  ammonia  and  the  caustic  alkalies, 
potassium  and  sodium  hydroxides ;  and  of  bleaching- 
powder  and  potassium  chlorate.  It  is  further  used  in 
large  quantity  in  the  process  of  tanning  for  the  pur- 
pose of  removing  the  hair  from  hides ;  in  decomposing 
fats  for  the  purpose  of  making  stearin  for  candles ; 
for  purifying  gas ;  and  especially  in  the  preparation  of 
mortar. 

Bleaching-powder. — The  preparation  of  bleaching-pow- 
der  was  referred  to  under  Chlorine  (which  see).  The 
main  reaction  involved  is  that  represented  in  the  equa- 
tion 

2Ca(OH)a  +  401  =  Ca(ClO)2  +  CaCl2  +  2H2O. 

Bleaching-powder 

The  compound  is  commonly  called  "  chloride  of  lime." 
Assuming  that  the  reaction  takes  place  in  the  same  way 
as  that  of  chlorine  on  caustic  potash,  the  product  is  a 
mixture  of  calcium  hypochlorite,  Ca(ClO)2,  and  calcium 
chloride,  for  it  is  held  that  the  reaction  with  potassium 
hydroxide  takes  place  as  represented  in  this  equation : 

2KOH  +  2C1  =  KC10  +  KC1  +  H2O. 

An  objection  to  the  view  that  calcium  chloride  is  pres- 
ent as  such  in  bleaching-powder  is  found  in  the  fact  that 
the  substance  is  not  deliquescent,  as  it  should  be  if  cal- 
cium chloride  were  present.  This  has  led  to  the  sug- 
gestion that  bleaching-powder  in  the  dry  form  is  not  a 
mixture  of  two  compounds,  as  represented  above,  but 

{Cl 
Od 

or  CaOCl2.  A  compound  of  this  formula  would  plainly 
have  the  same  composition  as  a  mixture  of  calcium  hy- 
pochlorite and  calcium  chloride  in  the  proportion  of 
their  molecular  weights.  For  we  have 


BLEACHING-POWDER.  535 

Ca(ClO)2  +  CaCl2  =  2CaOCl2. 

The  point  is  a  difficult  one  to  decide,  but  at  present  the 
evidence  appears  to  be  rather  in  favor  of  the  view  that 
bleaching-powder  in  the  dry  form  is  a  single  compound 
of  the  constitution  represented  by  the  last  formula 
given.  When  treated  with  water,  however,  it  appears  to 
be  resolved  into  a  mixture  of  the  hypochlorite  and  chlo- 
ride. 

Bleaching-powder  is  a  white  powder  which  has  the 
odor  of  hypochlorous  acid.  It  is  soluble  in  about 
twenty  parts  of  water,  though  the  commercial  product 
always  leaves  a  slight  residue,  which  consists  mainly  of 
calcium  hydroxide.  When  treated  with  an  acid,  as  sul- 
phuric or  hydrochloric  acid,  it  gives  up  all  its  chlorine. 
Thus,  with  hydrochloric  acid  the  reaction  takes  place 
as  represented  in  these  equations  : 

Ca(ClO)2  +  2HC1     =  CaCl2  +  2HC1O  ; 
2HC1  +  2HC10  =  2H2O  +  201,. 

With  sulphuric  acid  the  action  also  probably  takes  place 
in  two  stages.  The  acid  acts  upon  the  hypochlorite, 
setting  hypochlorous  acid  free ;  and  upon  the  chloride, 
setting  hydrochloric  acid  free.  The  hydrochloric  and 
hypochlorous  acids  then  react  with  each  other  as  repre- 
sented above : 

Ca(C10)2  +  H2S04   =  CaS04  +  2HC10  ; 
CaCl2       +H2S04  =CaSO4  +  2HCl; 
2HC1       +  2HC1O  =  2H2O   +  2C12. 

When  exposed  to  the  action  of  carbon  dioxide  hypo- 
chlorous  acid  is  liberated.  Hence,  when  it  is  allowed  to 
lie  in  the  air  this  decomposition  takes  place  slowly.  The 
hypochlorous  acid  acts  further  upon  the  calcium  chlo- 
ride, liberating  chlorine  : 

CaCl2  +  2HOC1  +  C02  =  CaCO3  +  H2O  +  201,. 

It  may  be,  however,  that  the  action  takes  place  between 
carbon  dioxide  and  the  compound  CaOCl2,  thus : 

CaOCl3  +  C02  =  CaCO3  +  Cla. 


536  INORGANIC  CHEMISTRY. 

In  any  case,  the  fact  remains  that  carbon  dioxide  sets 
the  chlorine  free  from  bleaching-powder. 

A  solution  of  bleaching-powder  alone  is  not  capable 
of  bleaching  except  very  slowly.  If,  however,  something 
is  added  which  has  the  power  to  decompose  it,  bleach- 
ing takes  place,  the  action  being  due  to  the  presence  of 
hypochlorous  acid  and  chlorine.  As  is  clear  from  what 
was  said  above,  the  passage  of  carbon  dioxide  through 
the  solution  or  the  addition  of  an  acid  would  cause  it  to 
bleach.  So,  too,  certain  salts  produce  a  similar  effect. 
The  explanation  of  this  is  the  instability  of  the  hypo- 
chlorites  formed  by  the  salts  added.  When  a  concen- 
trated solution  of  bleaching-powder  is  heated  it  gives  off 
oxygen,  and  the  salt  is  converted  into  the  chloride.  In 
dilute  solution,  however,  the  hypochlorite  is  converted 
into  chlorate  and  chloride  : 

3Ca(ClO)2  =  Ca(ClO3)2  +  2CaCl2. 

This  fact  is  taken  advantage  of,  as  has  been  shown,  for 
the  purpose  of  making  calcium  chlorate,  and  from  this 
potassium  chlorate  (see  p.  494).  In  contact  with  certain 
oxides,  as  copper  oxide,  ferric  oxide,  and  with  hydroxides, 
as  cobalt  and  nickel  hydroxides,  a  solution  of  bleaching- 
powder  readily  gives  up  oxygen  when  heated. 

The  chief  application  of  bleaching-powder  is,  as  its 
name  implies,  for  bleaching.  It  is  also  used  as  a  disin- 
fectant, and  as  an  antiseptic,  that  is,  for  the  purpose  of 
destroying  disease  germs,  and  of  preventing  decomposi- 
tion of  organic  substances. 

Calcium  Carbonate,  CaCO3. — This  salt  occurs  in  im- 
mense quantities  in  nature  in  the  well-known  forms  lime- 
stone, calc-spar,  marble,  and  chalk.  The  variety  of 
calc-spar  found  in  Iceland,  and  known  as  Iceland  spar, 
is  particularly  pure  calcium  carbonate.  It  crystallizes 
in  a  number  of  different  forms,  the  most  common  being 
in  rhombohedrons,  as  seen  in  ordinary  calc-spar.  A 
second  variety  of  crystallized  calcium  carbonate  is  ara- 
gonite.  This  is  found  in  nature  crystallized  in  rhombic 
prisms,  and  in  forms  derived  from  this.  When  heated 


CALCIUM  CARBONATE.  537 

aragonite  falls  to  pieces,  the  particles  being  small  crys- 
tals of  the  form  characteristic  of  calc-spar.  This  is  a 
case  of  dimorphism  similar  to  that  presented  by  sul- 
phur, which,  it  will  be  remembered,  crystallizes  in  two 
forms,  the  rhombic  and  monoclinic,  the  latter  of  which 
passes  into  the  former  spontaneously.  These  forms  are 
produced  artificially  very  readily.  When  calcium  car- 
bonate is  precipitated  from  a  solution  of  a  calcium  salt 
by  adding  a  soluble  carbonate  at  ordinary  temperatures 
the  precipitate  is  made  up  of  microscopic  crystals  which 
have  the  same  form  as  calc-spar.  If,  however,  the  solu- 
tion from  which  the  carbonate  is  precipitated  is  hot,  the 
salt  consists  of  microscopic  crystals  of  the  form  of  ara- 
gonite. 

The  most  abundant  form  of  calcium  carbonate  is  lime- 
stone, of  which  many  great  mountain-ranges  are  largely 
made  up.  This  is  a  compact  form  of  the  compound, 
which  has  a  gray  color,  and  frequently  consists  of  mi- 
nute crystals.  It  is  always  more  or  less  impure,  contain- 
ing clay  and  other  •  substances.  Limestone  which  is 
mixed  with  a  considerable  proportion  of  clay  is  called 
marl.  Many  natural  waters  contain  calcium  carbonate 
in  solution — probably  in  the  form  of  the  acid  carbonate. 
When  such  a  water  evaporates  the  carbonate  is  again 
deposited.  It  happens  in  some  places  that  a  water 
charged  with  the  carbonate  works  its  way  slowly  through 
the  earth  and  drops  from  the  top  of  a  cave.  Under  these 
circumstances  there  is  a  gradual  deposit  of  the  salt 
which  remains  suspended.  Such  hanging  formations  of 
the  carbonate  are  known  as  stalactites.  At  the  same 
time  that  part  of  the  liquid  which  falls  to  the  bottom  of 
the  cave  forms  a  projecting  mass  below  the  stalactite. 
Such  projecting  masses  are  called  stalagmites.  The  for- 
mation of  stalactites  takes  place  in  much  the  same  way  as 
that  of  icicles. 

Much  of  the  calcium  carbonate  found  in  nature  has  its 
origin  in  the  remains  of  animals,  and  fossils  are  very 
abundant  in  it.  Chalk  consists  almost  exclusively  of  the 
shells  of  microscopic  animals. 

When  carbon  dioxide  is  passed  into  a  solution  of  cal- 


538  INORGANIC  CHEMISTRY. 

cium  hydroxide,  the  carbonate  is  precipitated  ;  and,  it  the 
current  of  gas  is  continued  long  enough,  the  carbonate 
is  redissolved.  It  appears,  therefore,  that  calcium  car- 
bonate is  soluble  in  water  that  contains  carbonic  acid. 
It  is  probable  that  the  cause  of  this  is  to  be  found  in  the 
formation  of  an  acid  carbonate,  possibly  the  one  of  the 
formula  HO-OC-O-Ca-O-CO-OH.  No  positive  evi- 
dence, of  the  formation  of  this  substance  has,  however, 
been  furnished.  If  it  is  formed,  it  is  certainly  very  un- 
stable ;  for,  on  heating  the  solution  to  boiling,  the  normal 
carbonate  is  precipitated  and  carbon  dioxide  is  given  off. 
Natural  waters  which  come  in  contact  with  limestone 
gradually  tske  up  more  or  less  of  the  carbonate,  with  the 
aid  of  the  carbon  dioxide  of  the  air,  and  when  such  a 
water  is  boiled,  the  carbonate  is  thrown  down.  A  water 
containing  calcium  carbonate  in  solution  is  called  a  hard 
water;  and,  as  this  kind  of  hardness  is  easily  removed  by 
boiling,  it  is  called  temporary  hardness  in  order  to  dis- 
tinguish it  from  a  kind  which  is  not  removed  by  boiling, 
and  is  therefore  called  permanent  hardness.  Temporary 
hardness  is  further  removed  by  adding  lime  to  the  water, 
when  normal  carbonate  is  formed,  which  is  at  once  pre- 
cipitated. 

The  decomposition  of  calcium  carbonate  by  heat,  lead- 
ing to  the  formation  of  lime,  or  calcium  oxide,  and  carbon 
dioxide,  was  referred  to  on  p.  468. 

Applications. — Calcium  carbonate  is  used,  in  the  arts, 
for  a  great  many  purposes,  as  in  the  manufacture  of  glass ; 
as  a  flux  (see  p.  531)  in  many  important  metallurgical 
operations,  as  in  the  reduction  of  iron  from  its  ores ;  in 
the  preparation  of  lime  for  mortar ;  etc.  As  is  well  known, 
further,  marble  and  some  of  the  varieties  of  limestone 
are  extensively  used  in  building ;  and  large  quantities  of 
chalk  are  also  used. 

Calcium  Sulphate,  CaSO4- — This  compound  is  very 
abundant  in  nature.  The  principal  natural  variety  is 
gypsum,  which  occurs  in  crystals  containing  two  mole- 
cules of  water,  CaSO4  +  2H2O.  This  is  perhaps  derived 
directly  from  the  normal  acid  S(OH)6,  having  the  con- 


CALCIUM  SULPHATE.  539 

stitution  represented  by  the   formula  (HO)4S<Q>Oa. 

The  salt  of  the  formula  CaS04  also  occurs  in  nature,  and 
is  called  anhydrite.  A  granular  form  of  gypsum  is  called 
alabaster.  Calcium  sulphate  is  difficultly  soluble  in  hot 
and  cold  water,  but  its  solubility  is  markedly  increased 
by  the  presence  of  certain  other  salts ;  as,  for  example, 
sodium  chloride.  It  is  comparatively  easily  soluble  in 
hydrochloric  acid  and  in  nitric  acid.  When  heated  to 
100°,  or  a  little  above,  it  loses  nearly  all  its  water  and 
forms  a  powder  known  as  plaster  of  Paris,  which  has  the 
power  of  taking  up  water  and  forming  a  solid  substance. 
This  process  of  solidification  is  known  as  "  setting.'1 
Plaster  of  Paris  is  very  largely  used  in  making  casts,  on 
account  of  its  power  to  harden  after  having  been  made 
into  a  paste  with  water.  The  hardening  is  a  chemical 
process,  and  is  caused  by  the  combination  of  water  with 
the  salt  to  form  the  crystallized  variety : 

CaSO4  +  2H2O  =  (HO)4S<^>Ca. 

"When  heated  to  200°,  and  above,  all  the  water  is  given  off 
from  gypsum,  and  the  product  now  combines  with  water 
only  very  slowly,  and  is  of  no  value  for  making  casts.  In 
general,  the  higher  the  temperature  to  which  the  gypsum 
is  heated,  the  greater  the  difficulty  with  which  the  pro- 
duct combines  with  water. 

Many  natural  waters  contain  gypsum  in  solution.  Such 
waters  act  in  some  respects  like  those  which  contain  cal- 
cium carbonate.  With  soap,  for  example,  they  form  in- 
soluble compounds.  They  are  called  hard  waters. 
This  kind  of  hardness  is  not  removed  by  boiling,  and 
it  is  therefore  called  permanent  hardness.  Magnesium 
sulphate  acts  in  the  same  way,  producing  permanent 
hardness. 

When  calcium  sulphate  is  treated  with  a  solution  of  a 
soluble  carbonate,  it  is  decomposed,  forming  the  carbon- 
ate as  represented  in  the  equation 

CaS04  +  Na3C03  =  Na2S04  +  CaCO,. 


540  INORGANIC  CHEMISTRY. 

This  change  is  effected  simply  by  allowing  the  two  to 
stand  in  contact  at  the  ordinary  temperature. 

Besides  being  used  for  making  casts,  calcined  gypsum 
is  used  also  in  surgery  for  making  plaster-of-Paris  band- 
ages, and  as  a  fertilizer.  Its  action  as  a  fertilizer  is  be- 
lieved by  some  to  be  due  to  the  fact  that  it  has  the  power 
to  hold  ammonia  and  ammonium  carbonate  in  combina- 
tion, and  thus  to  make  them  available  for  the  plants.  It 
has  recently  been  shown  that  it  in  some  way  facilitates 
the  process  of  nitrification,  and  perhaps  it  is  in  conse- 
quence of  this  that  it  facilitates  plant-growth. 

Calcium  Phosphates. — There  are  three  phosphates  of 
calcium  :  (1)  The  normal  phosphate,  Ca3(PO4)2 ;  (2)  the 
secondary  phosphate,  CaHPO4 ;  and  (3)  the  primary  phos- 
phate, CaH4(PO4),. 

(1)  Normal  calcium  phosphate,  Ca3(PO4)2,  is  derived  from 
phosphoric  acid  by  the  replacement  of  all  the  hydrogen 
by  calcium.  It  is  found  in  nature  in  large  quantity  as 
phosphorite,  and  in  combination  with  calcium  fluoride  or 
chloride  as  apatite.  It  is,  further,  the  chief  inorganic 
constituent  of  bones,  forming  85  per  cent  of  bone-ash, 
and  is  contained  in  the  excrement  of  animals,  as  in  guano, 
etc.  It  is  found  everywhere  in  the  soil,  and  is  taken  up 
by  the  plants  for  whose  development  it  is  essential.  That 
it  is  also  essential  to  the  life  of  animals  is  obvious  from 
the  fact  that  the  bones  consist  so  largely  of  it.  The 
phosphate  needed  for  the  building  up  of  bones  is  taken 
into  the  system  with  the  food.  From  these  statements, 
it  is  clear  that  calcium  phosphate  is  of  fundamental 
importance,  and  that  a  fertile  soil  must  either  contain 
this  salt  or  something  from  which  it  can  be  formed. 
Now,  when  a  crop  is  raised  on  a  given  area,  a  certain 
amount  of  the  phosphate  contained  in  it  is  withdrawn. 
If  the  plants  were  allowed  to  decay  where  they  grow,  the 
phosphate  would  be  returned  and  the  soil  would  continue 
fertile  ;  but  in  cultivated  lands  this  is  not  the  case.  The 
crops  are  removed,  and  with  them  the  calcium  phos- 
phates contained  in  them,  and  the  soil  therefore  becomes 
exhausted.  If  the  substances  removed  are  used  as  food, 
some  of  the  phosphate  is  found  in  the  excrement  of  the 


CALCIUM  PHOSPHATES.  541 

animals  ;  and,  if  this  excrement  is  put  on  the  soil,  it  is 
again  rendered  fertile.  There  are,  however,  other  sources 
of  calcium  phosphate,  and  some  of  these  are  utilized  ex- 
tensively in  the  preparation  of  artificial  fertilizers.  The 
natural  form  of  the  phosphate,  as  that  in  bone-ash,  in 
phosphorite,  and  in  guano,  is  mainly  the  normal  or  neu- 
tral phosphate.  This  is  insoluble  in  water,  and  is  there- 
fore taken  up  by  the  plants  with  difficulty.  To  mate  it 
quickly  available,  it  must  be  converted  into  a  soluble 
phosphate.  This  is  done  by  treating  it  with  sulphuric 
acid  in  order  to  effect  the  reaction  represented  in  this 
equation  : 

Cas(PO4)2  +  2H2S04  =  CaH4(P04)2  +  2CaSO4. 

The  primary  phosphate  thus  formed  is  soluble  in  water, 
and  is  of  great  value  as  a  fertilizer.  The  mixture  of  the 
soluble  phosphate  and  of  calcium  sulphate  is  known  as 
"  superphosphate  of  lime."  The  sulphate,  as  we  have 
seen,  is  also  of  value  as  a  fertilizer.  The  value  of  super- 
phosphates depends  mostly  upon  the  amount  of  soluble 
phosphate  contained  in  them  ;  and  in  dealing  with  them 
it  is  customary  to  state  how  much  "  soluble  "  and  how 
much  "insoluble  phosphoric  acid"  they  contain.  When 
a  superphosphate  is  allowed  to  stand  for  a  time,  some  of 
the  soluble  primary  phosphate  is  converted  into  insol- 
uble phosphates  by  contact  with  basic  hydroxides  and 
water.  This  is  known  as  the  process  of  "reversion," 
and  that  part  of  the  phosphoric  acid  which  is  contained 
in  the  insoluble  phosphate  is  spoken  of  as  "  reverted 
phosphoric  acid." 

Normal  calcium  phosphate,  as  has  been  stated,  is  in- 
soluble in  water,  and  is  formed  when  a  soluble  normal 
phosphate  is  added  to  a  solution  of  a  calcium  salt.  It  is 
also  formed  when  di-sodium  phosphate  and  ammonia  are 
added  to  a  solution  of  a  calcium  salt,  thus  : 


+  3CaCl,  -f  2NH3  =  Ca3(PO4)2  +  4NaCl 

Di-sodium  phosphate  alone  at  first  produces  a  precipitate 
of  the  normal  phosphate,  while  the  primary  phosphate 
which  is  formed  at  the  same  time  remains  in  solution. 


542  INORGANIC  CHEMISTRY. 

The  reaction  takes  place  thus  : 

4HNa2P04  +  4CaCla  .=  CaH4(PO4)3  +  Ca3(PO4)3  +  SNaCL 

On  standing,  the  primary  acts  upon  the  tertiary  salt, 
forming  the  secondary  phosphate  thus  : 

CaH4(P04)2  +  Ca3(P04)2  =  4HCaPO4. 

But  even  on  long  standing  this  reaction  is  not  complete. 
Normal  or  tertiary  calcium  phosphate  is  soluble  in  hy- 
drochloric acid  and  in  nitric  acid,  in  consequence  of  the 
formation  of  calcium  chloride,  or  nitrate,  and  the  primary 
phosphate.  If  ammonia  is  added  to  this  solution,  the 
tertiary  phosphate  is  again  precipitated,  as  represented 
below  : 

Ca8(PO4)2  +  4HC1  =  2CaCl2  +  H4Ca(PO4)2 ; 
2CaCl2  +  H4Ca(P04)2  +  4NH8  =  Ca3(PO4)2  +  4NH4C1. 

(2)  Secondary  calcium  phosphate,  CaHPO4,  is  formed, 
as  above  described,  when  a  solution  of  a  calcium  salt 
is  treated  with  secondary  sodium  phosphate. 

(3)  Primary    calcium  phosphate,  H4Ca(PO4)2,  is  com- 
monly called  the  acid  phosphate  of  calcium.    It  is  formed 
when  ordinary  insoluble  calcium  phosphate  is  treated 
with  concentrated  sulphuric  acid,  and  is  contained  in  the 
so-called  superphosphates.     It  is  also  formed  by  treat- 
ing the  neutral  phosphate  with  phosphoric  acid  and  with 
hydrochloric  acid.     When  treated  with  but  little  water,  it 
is  converted  into  the  secondary  salt  and  free  acid : 

H4Ca(PO4)2  =HCaP04  +  H3PO4. 

Calcium  Silicate,  CaSiO3,  occurs  in  nature  as  the  mineral 
wollastonite,  and,  in  combination  with  other  silicates,  in 
a  large  number  of  minerals,  as  garnet,  mica,  the  zeolites, 
etc.  It  is  formed  when  a  solution  of  sodium  silicate  is 
added  to  a  solution  of  calcium  chloride,  and  when  a 
mixture  of  calcium  carbonate  and  quartz  is  heated  to  a 
high  temperature. 

Glass. — Ordinary  glass  is  a  silicate  of  calcium  and 
sodium  made  by  melting  together  sand  (silicon  dioxide, 
SiO2)  with  lime  and  sodium  carbonate  or  soda.  In- 


GLASS.  543 

stead  of  calcium  carbonate,  lead  oxide  may  be  used  ;  and 
instead  of  sodium  carbonate,  potassium  carbonate.  The 
properties  of  the  glass  are  dependent  upon  the  materials 
used  in  its  manufacture. 

Ordinary  window  glass  is  a  sodium-calcium  glass. 
The  purer  the  calcium  carbonate  and  silica,  the  better 
the  quality  of  the  glass.  This  glass  is  comparatively 
easily  acted  upon  by  chemical  substances,  and  is  there- 
fore not  adapted  to  the  preparation  of  vessels  which  are 
to  be  used  to  hold  acids  and  other  chemically  active 
substances.  It  answers,  however,  very  well  for  windows. 
The  difference  between  ordinary  window  glass  and  plate 
glass  is  essentially  that  the  former  is  blown  and  then 
cut  into  pieces,  while  the  latter,  when  in  the  molten  con- 
dition, is  run  into  flat  moulds  and  there  allowed  to  solidify. 

Bohemian  glass  is  made  with  potassium  carbonate.  If 
pure  carbonate  is  used,  as  well  as  pure  calcium  carbonate 
and  silica,  a  very  beautiful  glass  is  the  result.  It  is 
characterized  by  great  hardness,  by  its  difficult  fusibility, 
and  by  its  resistance  to  the  action  of  chemical  substances. 
It  is  particularly  well  adapted  to  the  manufacture  of 
vessels  and  tubes  for  use  in  chemical  laboratories. 

Flint-glass  is  made  by  melting  together  lead  oxide, 
potassium  carbonate,  and  silicon  dioxide.  It  is  charac- 
terized by  its  power  to  refract  light,  its  high  specific 
gravity,  its  low  melting-point,  and  the  ease  with  which 
it  is  acted  upon  by  reagents.  Owing  to  its  high  refrac- 
tive power,  it  is  largely  used  in  the  manufacture  of  lenses 
for  optical  instruments. 

Strass  is  a  variety  of  lead-glass  which  is  particularly 
rich  in  lead.  Its  refracting  power  is  so  great  that  it  is 
used  in  the  manufacture  of  artificial  gems. 

Colors  are  given  to  glass  by  putting  in  the  fused  mass 
small  quantities  of  various  substances.  Thus,  a  cobalt 
compound  makes  glass  blue  ;  copper  and  chromium  make 
it  green  ;  one  of  the  oxides  of  copper  makes  it  red  ;  ura- 
nium gives  it  a  yellow  color ;  etc.  The  most  common 
variety  of  glass  is  that  used  in  the  manufacture  of  ordi- 
nary bottles.  It  is  generally  green  to  black,  and  some- 
times brown.  In  its  manufacture,  impure  materials  are 


544  INORGANIC  CHEMISTRY. 

used,  chiefly  ordinary  sand,  limestone,  sodium  sulphate, 
common  salt,  clay,  etc. 

When  glass  is  suddenly  cooled,  it  is  very  brittle  and 
breaks  into  small  pieces  when  scratched  or  slightly 
broken  in  any  way.  This  is  shown  by  the  so-called 
Prince  Rupert's  drops,  which  are  made  by  dropping 
glass,  in  the  molten  condition,  into  water.  When  the  end 
of  such  a  drop  is  broken  off,  the  entire  mass  is  completely 
shattered  into  minute  pieces.  It  is  clear  from  this  that, 
in  the  manufacture  of  glass  objects,  care  must  be  taken 
not  to  cool  them  suddenly.  In  fact  they  are  cooled  very 
slowly,  the  process  being  known  as  annealing.  For  this- 
purpose  they  are  placed  in  furnaces  the  temperature  of 
which  is  but  little  below  that  of  fusion,  and  they  are 
kept  there  for  some  time,  the  heat  being  gradually 
lowered.  If  red-hot  glass  is  introduced  into  heated  oil 
or  paraffin,  and  allowed  to  cool  very  slowly,  it  is  found 
to  be  extremely  hard  and  elastic.  The  glass  of  De 
la  Bastie  is  made  in  this  way.  Vessels  made  of  it  can 
be  thrown  about  upon  hard  objects  without  breaking, 
but  sometimes  a  slight  scratch  will  cause  the  glass  to  fly 
in  pieces  as  the  Rupert's  drops  do. 

Mortar. — Mortar  is  made  of  slaked  lime  and  sand. 
When  this  mixture  is  exposed  to  the  air,  calcium  carbo- 
nate is  slowly  formed  and  the  mass  becomes  extremely 
hard.  The  water  contained  in  the  mortar  soon  passes 
off,  but  nevertheless  freshly  plastered  rooms  remain 
moist  for  a  considerable  time.  This  is  due  to  the  fact 
that  a  reaction  is  constantly  taking  place  between  the 
carbon  dioxide  and  calcium  hydroxide  by  which  calcium 
carbonate  and  water  are  formed, 

Ca(OH)2  +  CO2  =  CaCO3  +  H2O, 

and  it  is  the  water  thus  liberated  which  keeps  the  air 
moist.  The  complete  conversion  of  the  lime  into  car- 
bonate requires  a  very  long  time,  because  the  carbonate 
which  is  formed  on  the  surface  protects,  to  some  extent, 
the  lime  in  the  interior. 

It  is  generally  regarded  as  unhealthy  to  live  in  rooms 
with  freshly  plastered  walls,  because  the  air  is  constantly 


CALCIUM  CARBIDE.  545 

kept  moist  in  consequence  of  the  reaction  above  men- 
tioned. It  is,  however,  difficult  to  see  why  the  presence 
of  a  little  extra  moisture1  in  the  air  should  be  unhealthy  ; 
and,  if  there  is  any  danger  from  freshly  plastered  walls, 
it  seems  probable  that  the  cause  must  be  sought  for  else- 
where. It  is  possible  that  the  constant  presence  of 
moisture  in  the  pores  of  the  wall  interferes  with  the  im- 
portant process  of  diffusion,  and  that  therefore  when  the 
room  is  closed  this  natural  method  of  ventilation  cannot 
come  into  play. 

When  lime-stones  which  contain  magnesium  carbonate 
and  aluminium  silicate  in  considerable  quantities  are 
heated  for  the  preparation  of  lime,  the  product  does 
not  act  with  water  as  calcium  oxide  does,'  and  this  lime 
is  not  adapted  to  the  preparation  of  ordinary  mortar. 
On  the  other  hand,  it  gradually  becomes  solid,  in  con- 
tact with  water,  for  reasons  which  are  not  known.  Such 
substances  are  known  as  cements,  or  hydraulic  cements. 
Other  cements  besides  those  made  in  the  manner  men- 
tioned are  known. 

Calcium  Sulphide,  CaS,  is  formed  by  heating  calcium 
sulphate  with  charcoal.  It  is  remarkable  on  account 
of  the  fact  that  it  is  phosphorescent.  After  having 
been  exposed  to  sun-light,  it  continues  to  give  light  for 
some  time  afterward.  This  and  the  similar  compound, 
barium  sulphide,  are  now  used  quite  extensively  in  the 
preparation  of  luminous  objects,  such  as  match-boxes, 
clock-faces,  plates  for  house-numbers,  etc. 

Calcium  Nitride,  Ca3N2,  a  brown  mass,  is  formed  by 
heating  a  15-20  per  cent  amalgam  of  calcium  in  dry 
atmosphere  to  dull  redness.  It  is  somewhat  volatile, 
and  is  decomposed  by  water,  with  the  formation  of  cal- 
cium hydroxide  and  ammonia. 

Calcium  Carbide,  C2Ca,  is  formed  by  heating  lime  and 
carbon  together  in  an  electric  furnace,  when  the  reac- 
tion represented  by  the  following  equation  takes  place : 

CaO  +  30  =  C,Ca  +  CO. 

It  is  manufactured  on  the  large  scale,  the  form  of 
carbon  used  being  coke.  The  carbide  forms  a  crystal- 


546  INORGANIC  CHEMISTRY. 

line  mass.  That  of  average  quality  has  a  reddish  color. 
That  of  bad  quality  has  a  grayish  or  black  color. 
When  pure  it  is  colorless.  When  treated  with  water  it 
yields  acetylene,  CaH2 ,  a  gas  that  is  coming  into  practi- 
cal use  on  account  of  its  value  for  purposes  of  illumina- 
tion (see  Acetylene,  p.  374). 

STKONTIUM,  Sr  (At.  Wt.  86.95). 

Occurrence  and  Preparation. — Strontium  occurs  in 
nature  in  the  form  of  the  sulphate,  SrSO4,  as  celestite, 
and  in  the  form  of  the  carbonate,  SrCO3,  as  strontianite. 
The  latter  is  found  in  large  quantities  in  Westphalia. 
The  element  is  isolated  by  the  action  of  an  electric  cur- 
rent on  the  molten  chloride. 

Properties. — It  is  very  similar  to  calcium,  having  a 
metallic  lustre  and  a  brass-yellow  color.  It  is  oxidized 
by  contact  with  the  air,  and  decomposes  water  rapidly 
with  evolution  of  hydrogen,  which  does  not,  however, 
take  fire  spontaneously. 

Compounds  of  Strontium. — The  compounds  of  stron- 
tium are  very  similar  to  those  of  calcium.  Its  chloride 
has  not  the  same  attraction  for  water  that  calcium  chlo- 
ride has,  though  it  deliquesces  when  left  in  contact  with 
the  air.  The  oxide  is  not  easily  made  by  decomposition 
of  the  carbonate  by  heat,  as  the  carbonate  is  much  more 
stable  than  that  of  calcium.  It  is,  however,  prepared 
without  difficulty  by  heating  the  nitrate.  When  brought 
in  contact  with  water,  the  oxide  forms  the  hydroxide, 
which  is  analogous  to  calcium  hydroxide.  It  is  more 
easily  soluble  in  water  than  the  latter. 

Strontium  nitrate,  Sr(NO3)2,  is  made  in  considerable 
quantity  for  the  purpose  of  preparing  a  mixture  which, 
when  burned,  gives  a  red  light  (red-fire,  Bengal-fire).  It 
is  easily  made  by  dissolving  strontianite  or  strontium 
carbonate  in  nitric  acid. 

Strontium  sulphate,  SrSO4,  occurs  in  nature  in  beauti- 
ful crystals  as  the  mineral  celestite.  It  is  formed  when  a 
soluble  sulphate  is  added  to  a  solution  of  a  strontium 
salt.  In  solubility  it  lies  between  calcium  sulphate  and 
barium  sulphate. 


BARIUM. 


547 


BARIUM,  Ba  (At.  Wt.  136.39). 

Occurrence  and  Preparation. — Barium  occurs  in  nature 
in  the  same  forms  of  combination  as  strontium,  viz.,  as 
the  carbonate,  BaCO3,  in  witherite  ;  and  as  the  sulphate, 
BaSO4,  in  barite  or  heavy  spar.  It  is  prepared  by  elec- 
trolysis of  the  molten  chloride. 

Properties. — It  closely  resembles  calcium  and  stron- 
tium, being  a  yellow  metal,  which  is  oxidized  by  contact 
with  the  air  and  readily  decomposes  water  at  the  ordi- 
nary temperature. 

Barium  Chloride,  BaCl2  +  2H2O,  is  prepared  by  dissolv- 
ing barium  carbonate  in  hydrochloric  acid.  It  dissolves 
easily  in  water,  but  not  as  easily  as  the  chlorides  of 
strontium  and  calcium.  The  order  of  solubility,  begin- 
ning with  the  most  soluble,  is,  calcium,  strontium,  bar- 
ium,— the  same  as  in  the  case  of  the  sulphates. 

Barium  Hydroxide,  Ba(OH)2,  is  formed  by  dissolving 
barium  oxide  in  water,  just  as  calcium  hydroxide  is 
formed  by  treating  calcium  oxide  with  water.  In  hot 
water  it  is  much  more  easily  soluble  than  calcium  hy- 
droxide, and  it  is  also  more  easily  soluble  in  cold  water. 
As  such  a  solution  acts  in  the  same  general  way  as  lime- 
water,  it  is  frequently  used  in  the  laboratory  for  the 
purpose  of  detecting  carbon  dioxide,  barium  carbonate 
being  insoluble.  Like  lime-water,  it  has  an  alkaline 
reaction. 

Barium  Oxide,  BaO,  is  made  by  heating  the  nitrate, 
as  the  carbonate  is  not  easily  decomposed  by  heat.  The 
most  interesting  property  of  the  oxide  is  its  power  to 
take  up  oxygen  when  heated  to  a  dark  red  heat  in  the 
air  or  in  oxygen,  when  it  forms 

Barium  Peroxide  or  Dioxide,  BaO2. — This  peroxide  is  a 
white  powder  which  looks  like  the  simple  oxide.  When 
heated  to  a  temperature  a  little  higher  than  that  re- 
quired for  its  formation,  it  breaks  down  into  barium 
oxide  and  oxygen.  The  formation  of  the  peroxide  by 
heating  the  oxide  in  the  air,  and  the  decomposition  of 
the  peroxide  at  a  higher  heat,  make  it  possible  to  extract 
oxygen  from  the  air  and  to  obtain  it  in  the  free  state. 


548  INORGANIC  CHEMISTRY. 

This  method  of  preparing  oxygen  on  the  large  scale 
from  the  air  was  referred  to  under  Oxygen.  It  is  stated 
that  the  oxide  improves  with  use.  Specimens  which 
have  been  in  use  for  two  years  are  said  to  be  as 
efficient  as  at  first.  When  a  solution  of  hydrogen  di- 
oxide, H2O2,  is  added  to  a  solution  of  barium  hydroxide, 
a  precipitate  is  formed  which  has  the  composition  BaO9 
-f-  8H2O.  When  filtered  and  put  in  a  vacuum  over  sul- 
phuric acid,  it  loses  all  its  water  and  leaves  behind  pure 
dioxide.  The  dioxide  is  a  convenient  starting-point  in 
the  preparation  of  hydrogen  dioxide.  It  is  only  necessary 
to  treat  it  with  hydrochloric  acid  in  order  to  make  a  solu- 
tion of  hydrogen  dioxide.  The  solution  made  in  this 
way,  however,  contains  barium  chloride.  To  make  a 
solution  containing  nothing  but  the  dioxide,  pure  barium 
peroxide  is  treated  with  dilute  sulphuric  acid,  when  in- 
soluble barium  sulphate  is  formed  and  the  hydrogen 
dioxide  remains  in  solution  : 

BaO2  +  H2SO4  =  BaSO4  +  H2O2. 

It  is  interesting  to  compare  the  action  of  hydrochloric 
acid  on  barium  peroxide  and  on  the  corresponding  com- 
pound of  manganese.  With  the  latter,  as  we  have  seen, 
the  reaction  takes  place  as  represented  in  this  equation : 

MnOa  +  4HC1  =  MnCla  +  2HaO  +  01, ; 

while  with  barium  peroxide  the  reaction  takes  place 
thus: 

BaOa  +  2HC1  =  BaCl3  +  H3Oa- 

It  is  probable  that  in  the  case  of  manganese  dioxide 
some  intermediate  reactions  take  place  which  are  im- 
possible in  the  other  case.  (See  Manganese  Dioxide.) 

Barium  Sulphide,  BaS,  is  made  as  calcium  sulphide  is, 
by  reducing  the  sulphate  by  heating  with  charcoal.  It 
is  phosphorescent,  like  the  calcium  compound.  When 
dissolved  in  water,  it  is  decomposed,  forming  the  hydro- 
sulphide  and  hydroxide,  thus : 


COMPOUNDS  OF  BARIUM.  549 

2BaS  +  2H20  =  Ba(SH)2  +  Ba(OH)2. 

It  will  be  remembered  that  thermochemical  investiga- 
tions have  made  it  appear  probable  that  similar  reac- 
tions take  place  when  potassium  and  sodium  sulphides 
are  dissolved  in  water.  In  the  case  of  barium  sulphide 
the  evidence  is  more  tangible,  for,  on  evaporating  a  so- 
lution of  this  compound,  both  the  hydrosulphide  and 
hydroxide  crystallize  out. 

Barium  Nitrate,  Ba(NO3)2,  is  easily  soluble  in  water, 
but  difficultly  soluble  in  acids,  and  is  precipitated  from 
its  solution  in  water  by  the  addition  of  nitric  acid. 
When  heated  to  a  sufficiently  high  temperature,  it  is  de- 
composed, and  barium  oxide  is  left  behind. 

Barium  Sulphate,  BaSO4. — This  occurs  in  nature  as 
barite,  or  heavy  spar,  and  is  precipitated  when  a  soluble 
sulphate  or  sulphuric  acid  is  added  to  a  solution  of  a 
barium  salt.  It  is  insoluble  in  water  ;  when  freshly  pre- 
cipitated, it  is  easily  soluble  in  concentrated  sulphuric 
acid.  It  is  artificially  prepared  for  use  as  a  pigment 
and  is  known  as  permanent  ivhite.  On  account  of  its  in- 
solubility it  is  much  used  in  chemical  analysis  for  the 
purpose  of  detecting  and  estimating  sulphuric  acid.  It 
differs  markedly  from  calcium  and  strontium  sulphate,  in 
the  fact  that,  when  treated  with  a  solution  of  ammonium 
carbonate,  it  is  not  converted  into  the  carbonate,  whereas 
calcium  and  strontium  sulphates  are  by  this  means  com- 
pletely converted  into  the  carbonates.  This  fact  is  taken 
advantage  of  in  analysis.  There  are  other  differences, 
which  will  be  stated  at  the  «end  of  this  chapter. 

Barium  Carbonate,  BaCO3,  occurs  in  nature  as  witherite, 
and  is  made  pure  by  adding  ammonium  carbonate  and  a 
little  ammonia  to  a  solution  of  barium  chloride.  The 
carbonate  usually  found  in  the  market  is  made  by  pre- 
cipitating a  solution  of  the  crude  sulphide  with  sodium 
carbonate,  or  by  heating  together  sodium  carbonate  and 
natural  barium  sulphate,  or  heavy  spar.  Made  in  either 
of  these  ways  it  contains  alkaline  carbonate,  from  which 
it  is  impossible  to  separate  it  by  washing.  The  carbonate, 
like  the  other  salts  of  barium,  is  poisonous.  It  has  the 
power  to  unite,  and  form  insoluble  compounds,  with  me- 


550  INORGANIC  CHEMISTRY. 

tallic  oxides  of  the  formula  M2O3,  as,  for  example,  ferric 
oxide,  Fe2O3,  and  is  used  in  analytical  operations  for  the 
purpose  of  separating  iron  from  other  metals,  like  man- 
ganese, which  are  not  precipitated  by  it. 

Phosphates  of  Barium.  —  The  phosphates  of  barium  cor- 
respond in  general  to  those  of  calcium.  When  ordinary 
sodium  phosphate  and  ammonia  are  added  to  a  solution 
of  a  barium  salt,  normal  or  tertiary  phosphate  is  pre- 
cipitated : 


3BaCla  +  2HNa3PO4  4-  2NH3  =  Ba3(PO4)2  +  4NaCl  +  2NH4C1. 

When  sodium  phosphate  alone  is  added,  the  first  reac- 
tion which  takes  place  is  that  represented  in  the  equa- 
tion 

4BaCl,  +  4HNaaP04  =  BaH4(PO4)2  +  Ba3(P04),  +  SNaCl. 

The  precipitate  is  the  tertiary  phosphate,  while  the 
primary  phosphate  is  in  solution.  On  standing,  the  solu- 
ble salt  acts  upon  the  insoluble  one,  forming  the  second- 
ary phosphate  thus  : 

BaH4(PO,),  +  Ba3(P04),  =  4HBaPO, 

Reactions  which  are  of  Special  Value  in  Analysis.  —  The 
sulphates  of  calcium  and  strontium  are  completely  con- 
verted into  the  carbonates  by  contact  with  a  solution 
of  ammonium  carbonate  in  ammonia.  The  sulphate  of 
barium  is  not  changed  in  this  way.  Consequently,  if  a 
mixture  of  the  three  sulphates  is  treated  with  ammoni- 
um carbonate,  those  of  calcium  and  strontium  will  be 
converted  into  carbonates,  while  that  of  barium  will  re- 
main unchanged.  By  filtering,  washing  with  water,  and 
treating  with  dilute  nitric  or  hydrochloric  acid,  the  car- 
bonates will  be  dissolved,  while  the  sulphate  will  not. 
If  nitric  acid  is  used,  the  solution  may  be  evaporated  to 
dryness  and  treated  with  a  mixture  of  alcohol  and  ether. 
Calcium  nitrate  will  dissolve  ;  strontium  nitrate  will  not. 

Fluosilicic  acid  produces  a  precipitate  of  barium  fluo- 
silicate,  BaSiFa,  in  solutions  of  barium  salts.  This  is 


GLUCINUM.  551 

insoluble  in  a  mixture  of  alcohol  and  water,  and  difficultly 
soluble  in  water.  The  corresponding  salts  of  calcium 
and  strontium  are  soluble. 

Calcium  sulphate  solution  produces  a  precipitate  in  a 
solution  of  a  strontium  salt  or  of  a  barium  salt,  but  not 
in  one  of  a  calcium  salt. 

Strontium  sulphate  solution  precipitates  barium  sul- 
phate from  a  solution  of  a  barium  salt,  but  forms  no  pre- 
cipitate in  a  solution  of  a  strontium  salt. 

When  boiled  with  a  solution  of  one  part  of  sodium 
carbonate  and  three  parts  of  sodium  sulphate,  the  sul- 
phates of  strontium  and  calcium  are  completely  con- 
verted into  carbonates,  while  the  sulphate  of  barium 
remains  unchanged. 

Barium  chloride  is  insoluble  in  absolute  alcohol ;  cal- 
cium chloride  is  easily  soluble ;  while  strontium  chloride 
dissolves  in  warm  absolute  alcohol. 

Ammonium  oxalate,  (NH4)2C204,  produces  precipitates 
of  the  oxalates  in  solutions  of  calcium,  barium,  and 
strontium.  Only  the  calcium  salt  is  insoluble  in  dilute 
acetic  acid. 

Potassium  dichromate,  K2Cr2O7,  precipitates  barium 
chromate,  BaCrO4.  The  corresponding  salts  of  calcium 
and  strontium  are  soluble  in  water.  Barium  chromate 
is  easily  soluble  in  hydrochloric  or  nitric  acid. 

All  three  elements  of  the  group  give  colored  flames 
which  have  characteristic  spectra.  Calcium  compounds 
color  the  flame  reddish  yellow ;  strontium  compounds  give 
an  intense  red  ;  and  barium  compounds  a  yellowish  green 
color.  The  spectra  are  more  complicated  than  those  of  the 
elements  of  the  potassium  group,  but  each  one  contains 
highly  characteristic  lines  which  are  easily  recognized. 

MAGNESIUM  SUB-GROUP. 
GLUCINUM,  Gl  (At.  Wt.  9.01). 

Occurrence  and  Preparation — The  principal  form  in 
which  the  element  glucinum  occurs  in  nature  is  in  the 
mineral  beryl,  which  is  a  silicate  of  aluminium  and  glu- 
cinum of  the  formula  Al2Gl3(SiO3)6.  Emerald  has  the 


552  INORGANIC  CHEMISTRY. 

same  composition,  but  is  colored  green  by  the  presence 
of  a  little  chromic  oxide.  The  element  can  be  isolated 
by  decomposing  the  chloride  by  heating  it  with  potas- 
sium or  sodium. 

Properties. — The  statements  concerning  the  properties 
of  glucinum,  made  by  those  who  have  prepared  it  in 
different  ways,  differ  somewhat  from  one  another,  evi- 
dently in  consequence  of  the  fact  that  it  has  not  gener- 
ally been  pure.  It  has  a  metallic  lustre.  When  heated 
in  the  flame  of  the  blow-pipe  it  becomes  covered  with  a 
thin  layer  of  oxide,  which  prevents  further  action ;  it 
dissolves  readily  in  hydrochloric  and  sulphuric  acids, 
but  only  with  difficulty  in  nitric  acid.  It  is  dissolved 
by  potassium  hydroxide,  forming  in  all  probability  a 
glucinate  of  the  composition  Gl(OK)a : 

Gl  +  2KOH  =  Gl(OK)a  +  H2. 

The  specific  heat  of  glucinum  at  ordinary  temperature 
is  0.425.  This  multiplied  by  the  atomic  weight  9.01  gives 
3.83  instead  of  6.24.  But  the  analysis  and  the  determina- 
tion of  the  specific  gravity  of  the  vapor  of  the  chloride 
show  that  it  has  the  formula  G1C1,,  the  atomic  weight  of 
glucinum  being  9.01.  At  257°  C.,  however,  the  specific 
heat  of  glucinum  is  0.582,  and  this  multiplied  by  9.01 
gives  5.24.  At  ordinary  temperature,  therefore,  glu- 
cinum, like  carbon,  boron,  and  silicon,  is  an  exception  to 
the  law  of  Dulong  and  Petit,  while  at  a  higher  tempera- 
ture, like  the  elements  named,  it  conforms  to  the  law. 

Compounds  of  Glucinum.  —  The  compounds  of  glu- 
cinum differ  in  many  respects  from  those  of  the  group 
calcium,  barium,  strontium.  The  hydroxide  is  entirely 
insoluble  in  water ;  the  sulphate  is  easily  soluble  in 
water ;  the  chloride  is  completely  decomposed  when  its 
water  solution  is  evaporated  to  dryness,  the  products 
being  hydrochloric  acid  and  glucinum  oxide.  It  shows 
a  marked  tendency  to  form  basic  salts. 

Glucinum  Chloride,  G1C12 ,  is  formed  by  the  action  of 
chlorine  on  glucinum,  and  more  easily  by  treating  a  mix- 
ture of  glucinum  oxide  and  carbon  with  chlorine,  the 
reaction  being  similar  to  that  employed  in  making  sili- 


COMPOUNDS  OF  OLUCINUM.  553 

con  chloride  (see  p.  413)  and  boron  chloride  (see  p.  352). 
It  is  volatile,  and  it  is  therefore  possible  to  determine 
the  specific  gravity  of  its  vapor.  This  has  been  done, 
with  the  result  of  showing  its  molecular  weight  to  be 
79.9.  Taking  this  fact  into  consideration,  together  with 
the  percentage  composition  of  the  compound,  the  con- 
clusion is  justified  that  the  atomic  weight  of  glucinum  is 
9.01.  For  a  long  time  it  was  thought  to  be  13.65,  with 
which  figure  the  specific  heat,  0.425,  is  in  accordance ; 
for  13.65  X  0.425  =  5.79,  but  the  evidence  furnished  by 
the  specific  gravity  of  the  vapor  of  the  chloride  is  re- 
garded as  conclusive  in  favor  of  the  atomic  weight  9.01. 

Glucinum  Hydroxide,  G1(OH)2 ,  is  thrown  down  as  a 
precipitate  when  a  soluble  hydroxide  is  added  to  a  solu- 
tion of  a  glucinum  salt : 

G1SO4  +  2NaOH  =  Gl(OH),  +  Na2SO4. 

It  is  a  white,  gelatinous  mass,  which  is  soluble  in  potas- 
sium and  sodium  hydroxides  and  in  ammonia,  so  that, 
after  precipitation  from  glucinum  salts  by  these  reagents, 
it  redissolves.  This  solution  is  due  to  the  formation  of 
glucinates  of  the  formula  G1(OM)2 : 

Gl(OH),  +  2NaOH  =  Gl(ONa)2  +  2H2O. 

When  sufficiently  diluted  with  water,  the  potassium  and 
sodium  salts  are  completely  decomposed,  and  the  hy- 
droxide reprecipitated.  This  is  an  illustration  of  mass 
action,  a  large  quantity  of  water  effecting  a  decomposition 
which  a  small  quantity  does  not  effect.  The  power  of 
the  hydroxide  to  form  salts  with  the  strong  bases  shows 
that  it  has  slight  acid  properties.  The  hydroxides  of  cal- 
cium, barium,  and  strontium  do  not  possess  this  power. 
Glucinum  Sulphate,  G1SO4,  is  formed  by  dissolving 
glucinum  hydroxide  in  dilute  sulphuric  acid,  and  has 
the  composition  G1SO4  -)-  4H2O  when  crystallized  from 
water.  When  a  solution  of  this  salt  is  heated  with  glu- 
cinum hydroxide,  basic  salts  are  formed,  of  which  the 
following  are  examples :  G12SO5  and  G13SO6.  The  first 
of  these  is  to  be  regarded  as  derived  from  a  hydroxide 
of  the  formula  HO-G1-O-G1-OH,  by  neutralization 


554  INORGANIC  CHEMISTRY. 

with    sulphuric    acid,    as    represented    in   the    formula 

r^i—  o 

O<Qi_Q>SO2;   the  second   from   a  hydroxide  of   the 

formula  HO-G1-O-G1-O-G1-OH,  by  neutralization  with 
sulphuric  acid,  as  represented  in  the  formula 

0    Gl-O  ) 
"<G1      I  SO, 

<G1-OJ 

With  potassium  sulphate,  glucinum  sulphate  forms  a 
double  salt  of  the  formula  K5SO,.  G1SO,  +  2H,O, 


,  (HO),SO<°K 

or  X>CH,     or     possibly  X>G1. 


Glucinum  Carbonate,  G-1CO3.  —  When  a  slight  excess  of 
sodium  carbonate  is  added  to  a  solution  of  glucinum 
sulphate,  a  basic  carbonate  of  the  formula  G13CO5  is 
formed.  This  is  similar  to  the  second  of  the  above-men- 
tioned basic  sulphates.  It  is  to  be  regarded  as  derived 
from  the  hydroxide  HO-G1-O-G1-O-G1-OH  by  neu- 
tralization with  carbonic  acid,  as  represented  in  the 


formula  XG1          CO. 


Weak  Basic  Character  of  Glucinum.  —  The  power  of 
glucinum  hydroxide  to  form  salts  with  strong  bases, 
such  as  potassium  and  sodium  hydroxides,  which  was  re- 
ferred to  above,  shows  that  the  hydroxide  has  slight  acid 
properties.  At  the  same  time,  as  we  should  expect,  its 
basic  properties  are  weaker  than  those  of  the  other  base- 
forming  elements  thus  far  considered.  This  is  shown  in 
the  ready  formation  of  basic  salts,  such  as  the  basic  sul- 
phates and  basic  carbonates  mentioned.  The  strongest 
bases  do  not  readily  form  basic  salts,  but  are,  on  the  other 
hand,  more  competent  to  form  stable  acid  salts.  Thus, 
potassium  and  sodium  form  acid  carbonates  ;  calcium 
appears  to  form  an  extremely  unstable  acid  carbonate, 
but  preferably  all  the  members  of  the  calcium  group 
form  normal  carbonates  of  the  general  formula  MCO3  ; 


MAGNESIUM.  555 

glucinum,  however,  and,  as  we  shall  see,  magnesium, 
preferably  form  basic  carbonates.  We  shall  see,  further, 
that  the  members  of  the  next  family,  of  which  aluminium 
is  the  principal  one,  form  only  extremely  unstable  com- 
pounds with  carbonic  acid,  their  basic  properties  not 
being  sufficiently  strong  to  hold  them  in  combination 
with  the  weak  acid,  except  apparently  at  a  very  low 
temperature. 

This  resemblance  to  the  acid-forming  elements  is 
shown  by  glucinum  also  by  the  ease  with  which  its 
chloride  is  decomposed  into  the  oxide  and  hydrochloric 
acid  when  its  water  solution  is  evaporated  to  dryness. 
This  reaction  does  not  take  place  in  the  case  of  sodium 
and  potassium  at  all,  nor  with  barium  and  strontium. 
With  calcium  it  takes  place  to  a  slight  extent,  but  with 
glucinum  it  is  complete,  as  it  is  with  the  similar  metal 
magnesium.  In  general,  the  more  acidic  the  element 
the  more  easily  is  its  chloride  decomposed  in  this  way. 

MAGNESIUM,  Mg  (At.  Wt.  24.10). 

Occurrence. — Magnesium  occurs  very  abundantly  in 
nature,  though  by  no  means  as  abundantly  as  calcium. 
Among  the  widely  distributed  minerals  which  contain 
the  element  are  magnesite,  which  is  the  carbonate, 
MgCO3 ;  dolomite,  a  double  carbonate  of  magnesium 
and  calcium ;  serpentine,  talc,  soapstone,  meerschaum, 
hornblende,  all  of  which  contain  magnesium  silicates. 
Further,  the  metal  is  found  in  solution  in  many  spring- 
waters  in  the  form  of  the  sulphate,  or,  as  it  is  called, 
Epsom  salt.  Kainite  is  a  sulphate  and  chloride  of  the 
composition  expressed  by  the  formula 

K2S04.MgS04.MgCl2  +  6H,0 ; 

kieserite  is  magnesium  sulphate,  MgSO4  -\-  H2O ;  car- 
nallite  is  a  double  chloride,  KMgCl3  +  6H2O. 

Magnesium  compounds  are  contained  in  the  soil  in 
consequence  of  the  decomposition  of  minerals  contain- 
ing it.  It  is  to  some  extent  taken  up  by  the  plants,  and 


556  INORGANIC  CHEMISTRY. 

subsequently  into  the  animal  body.     It  is  found  in  the 
bones  and  in  the  blood  in  small  quantities. 

Preparation. — The  metal  can  be  made  by  the  electrol- 
ysis of  its  chloride,  but  is  most  conveniently  made  by 
decomposing  the  chloride  by  means  of  sodium.  It  is 
now  manufactured  in  considerable  quantity  by  this 
method.  The  operation  consists  in  bringing  together 
dry  magnesium  chloride,  fluor-spar,  and  sodium  in  cer- 
tain proportions,  and  heating  to  a  high  temperature  in  a 
crucible.  The  metal  is  purified  by  distillation.  Instead 
of  using  the  chloride,  which  it  is  difficult  to  prepare  dry 
in  large  quantity,  the  double  chloride  of  magnesium  and 
potassium,  KMgCl3  or  MgCl2.KCl,  is  frequently  used. 

Properties. — It  is  a  silver-white  metal  with  a  high 
lustre.  In  the  air  it  changes  only  slowly,  but  it  gradu- 
ally becomes  covered  with  a  layer  of  the  hydroxide.  At 
ordinary  temperatures  magnesium  does  not  decompose 
water ;  at  100°  it  decomposes  it  slowly.  When  heated 
above  its  melting-point  in  oxygen  or  in  the  air,  it  takes 
fire  and  burns  with  a  bright  flame,  forming  the  white 
oxide.  The  light  of  the  flame  is  very  efficient  in  produc- 
ing certain  chemical  changes,  such  as  those  involved  in 
photography,  when  a  permanent  impression  is  made  by 
the  light  upon  a  sensitive  plate.  It  has  also  the  power 
to  cause  hydrogen  and  chlorine  to  combine  just  as  the 
sunlight  and  the  electric  light  do. 

Applications. — The  principal  use  to  which  magnesium 
is  put  is  for  the  purpose  of  producing  a  bright  light,  as  for 
photographing  in  spaces  to  which  the  sunlight  does  not 
have  access,  and  for  signaling.  It  is  also  used  to  some 
extent  as  an  ingredient  of  materials  employed  in  making 
fireworks. 

Compounds  of  Magnesium. — The  compounds  of  mag- 
nesium present  a  general  resemblance  to  those  of  gluci- 
iium.  As  the  element  is  much  more  abundant  in  nature, 
its  compounds  have  been  studied  more  extensively.  Its 
acid  properties  are  somewhat  weaker,  and  its  basic 
properties  stronger,  than  those  of  glucinum.  Its  hy- 
droxide does  not  form  salts  with  the  hydroxides  of 
potassium  and  sodium.  On  the  other  hand,  its  chlo- 


MAGNESIUM  CHLORIDE.  557 

ride  is  decomposed  when  its  water  solution  is  evap- 
orated to  dryness.  The  hydroxide  is  very  slightly  sol- 
uble in  water,  and  this  solution  has  a  slightly  alkaline 
reaction.  With  carbonic  acid  it  forms  basic  carbon- 
ates similar  to  those  formed  by  glucinum.  On  the 
other  hand,  it  does  not  readily  form  basic  salts  with 
sulphuric  acid.  In  character,  it  is  plainly  more  closely 
allied  to  the  members  of  the  calcium  group  than  gluci- 
num is. 

Magnesium  Chloride,  MgCl2. — This  salt,  as  has  been 
stated,  occurs  in  nature.  It  is  easily  formed  by  dissolv- 
ing magnesium  oxide  or  carbonate  in  hydrochloric  acid. 
On  evaporating  at  as  low  a  temperature  as  possible, 
there  finally  crystallizes  out  of  the  very  concentrated 
solution,  a  salt  of  the  composition  MgCl2  -f-  6H2O,  anal- 
ogous to  crystallized  calcium  chloride,  CaCl2  +  6H2O, 
and  strontium  chloride,  SrCl2  -f-  6H2O.  When  this 
crystallized  salt  is  heated  for  the  purpose  of  driving 
off  the  water,  it  is  completely  decomposed  in  accordance 
with  the  following  equation  : 

MgCl2  +  H20  =  MgO  +  2HC1. 

It  is  most  conveniently  prepared  in  the  dry  form  by 
first  making  ammonium-magnesium  chloride,  and  de- 
composing this  by  heat.  For  this  purpose,  a  solution  of 
ammonium  chloride  is  added  to  a  solution  of  magnesium 
chloride  and  the  whole  evaporated  to  dryness.  There  is 
formed  in  the  solution  the  double  salt  of  the  composi- 
tion NH4MgCl3(MgCl2.NH4Cl),  which  can  be  evaporated 
to  complete  dryness.  When  perfectly  dry,  this  double 
salt  breaks  down  into  magnesium  chloride  and  ammo- 
nium chloride,  if  heated  to  a  sufficiently  high  tempera- 
ture. The  ammonium  chloride  under  these  circum- 
stances is  volatilized,  and  the  magnesium  chloride  re- 
mains behind  in  the  molten  condition. 

The  chloride  is  a  white,  crystalline  mass  which  de- 
liquesces in  the  air.  At  a  bright  red  heat,  it  is  volatile 
and  can  be  distilled  in  an  atmosphere  of  hydrogen.  It 
dissolves  in  water  with  marked  evolution  of  heat.  It 
-combines  readily  with  the  chlorides  of  potassium,  so- 


558  INORGANIC  CHEMISTRY. 

dium,  and  ammonium,  forming  crystallizing  compounds 
of  the  formulas  KMgCl3,  NaMgCl,,  and  NH.MgCl,,  which 
may  be  regarded  as  formed  by  the  combination  of  one 
molecule  of  magnesium  chloride  with  one  molecule  of  each 
of  the  other  chlorides.  A  second  compound  with  potas- 
sium chloride,  of  the  formula  K2MgCl4,  is  also  known. 
It  seems  probable  that  the  latter  is  analogous  to  the  po- 

OK 

tassium  compound  of  glucinum  of  the  formula  GI 


It  corresponds  to  an  oxygen  compound  of  the  formula, 

OK 

,  which,  however,  does  not  seem  to  be  formed. 


If  in  this  compound  we  imagine  each  of  the  two  oxygen 
atoms  to  be  replaced  by  two  chlorine  atoms,  the  com- 

//^l  \_T7" 

pound  would  have  the  formula  Mg  <  >QI  2<_g-  .    The  exist- 

ence of  two  double  chlorides  of  magnesium  and  potas- 
sium is  suggested  by  what  has  been  said  regarding 
compounds  of  this  kind  (see  p.  465).  One  of  these 

Cl 
would  be  represented  by  the  formula  Mg  <  /  ^  -,  yrr,    the 

other  by  the  above  formula.  Both  are  known. 
Further,  magnesium  bromide  forms  the  salt  K2MgBrt 


or 

Magnesium  Oxide,  MgO.  —  This  compound  is  commonly 
called  magnesia.  A  fine  white  variety  which  is  known 
as  magnesia  usta,  is  made  by  heating  precipitated  basic 
magnesium  carbonate.  It  is  a  white,  loose  powder, 
which  is  very  difficultly  soluble  in  water,  forming  with  it 
the  hydroxide,  Mg(OH)2,  which  is  also  very  difficultly 
soluble.  Magnesia  is  used,  in  medicine,  as  an  applica- 
tion to  wounds,  and,  mixed  with  a  solution  of  ferric  sul- 
phate, as  an  antidote  in  cases  of  poisoning  by  arsenic. 
As  magnesia  is  infusible,  it  is  used  to  protect  vessels 
which  are  subjected  to  a  high  temperature.  When 
mixed  with  water  and  allowed  to  lie  in  the  air,  it  be- 
comes very  hard.  Mixtures  of  magnesia  with  sand  also 
have  this  property,  and  are  used  as  hydraulic  cements. 
It  is  used,  further,  in  the  manufacture  of  fire-bricks. 


MAGNESIUM  CARBONATE.  559 

Magnesium  Sulphate,  MgSO4.  —  The  mineral  kieserite, 
which  occurs  in  Stassfurt,  has  the  composition  expressed 
by  the  formula  MgSO4  +  H2O  ;  or,  more  probably,  this 

oH 


should  be  written  (HO)2MgSO3,  or  OS  -  ,  in  which 


it  appears  as  a  derivative  of  the  acid  SO(OH)4.  The 
salt  MgSO4  +  7H2O  (or  H2MgSOB  +  6H2O),  also  occurs 
in  nature.  It  is  this  variety  which  is  generally  obtained 
when  a  solution  of  magnesium  sulphate  is  evaporated  to 
crystallization.  It  crystallizes  in  large  rhombic  prisms, 
or,  if  rapidly  deposited  from  very  concentrated  solutions, 
in  small,  needle-shaped  crystals.  At  ordinary  tempera- 
tures, 100  parts  of  water  dissolve  125  parts  of  the  salt. 
The  water  solution  has  a  bitter,  salty  taste.  When 
heated,  it  readily  loses  6  molecules  of  water,  but  it  re- 
quires a  temperature  of  over  200°  to  drive  off  the  last 
molecule.  This  has  led  to  the  belief  that  the  salt  with 
one  molecule  has  the  constitution  above  given,  being  a 
derivative  of  the  acid  SO(OH)4. 

Magnesium  sulphate  finds  extensive  application.  It  is 
used  in  medicine  as  a  purgative,  and  is  known  as  Ep- 
som salt,  for  the  reason  that  it  is  contained  in  the  water 
of  Epsom  springs  ;  it  is  used  further  in  the  manufac- 
ture of  sodium  sulphate  and  potassium  sulphate,  and  as 
a  fertilizer  in  place  of  gypsum,  it  having  been  shown  to 
be  advantageous  in  some  cases.  Its  chief  use  is  for 
loading  cotton  fabrics. 

Magnesium  sulphate  forms  double  salts  with  other 
sulphates  ;  as,  for  example,  one  with  potassium  sulphate, 
similar  to  that  formed  by  glucinum  sulphate  (see  p.  554). 
The  constitution  of  the  double  sulphate  of  magnesium 
and  potassium  is  probably  that  expressed  by  the  for- 

o,s<gK 

mula  n  >  Mg. 


Magnesium  Carbonate,  MgCO3.  —  Like  glucinum,  mag- 
nesium shows  a  marked  tendency  to  form  basic  salts 


560  INORGANIC  CHEMISTRY. 

with  carbonic  acid.  When  a  neutral  magnesium  salt  is 
treated  with  a  soluble  carbonate,  a  basic  carbonate  is 
precipitated,  the  composition  of  which  varies  according 
to  the  conditions  under  which  it  is  prepared.  The  salt 
obtained  by  adding  an  excess  of  sodium  carbonate  to  a 
solution  of  magnesium  sulphate  has  the  composition 
Mg3(OH)2(CO3)2  It  is  derived  from  three  molecules 
of  magnesium  hydroxide  and  two  of  carbonic  acid,  as  is 

00<0-Mg-OH 

more  clearly  shown  in  the  formula  n>Mg       .    The 

c*c\  *>** 

<O-Mg-OH 

salt  which  is  manufactured  on  the  large  scale  is  more 
complicated  than  this,  being  derived  from  four  molecules 
of  magnesium  hydroxide  and  three  of  carbonic  acid.  It 
is  known  as  magnesia  alba.  It  is  this  form  of  the  car- 
bonate which  is  used  in  the  preparation  of  magnesia  usta. 
Normal  magnesium  carbonate,  MgCO3,  occurs  in  nature 
as  magnesite.  It  crystallizes  in  the  same  form  as  cal- 
cium carbonate,  or  is  isomorphous  with  it.  It  is  insolu- 
ble in  water,  but  like  calcium  carbonate  it  dissolves  in 
water  containing  carbon  dioxide  in  solution.  From  this 
solution  crystals  having  the  composition  MgCO3  -f~~ 
3H2O  and  MgCO3  -f-  5H2O  are  deposited  under  the 
proper  conditions. 

Phosphates. — The  conduct  of  the  phosphates  of  mag- 
nesium is  very  similar  to  that  of  the  phosphates  of  cal- 
cium. All  three  are  known  ;  and  of  these  only  the  primary 
salt  is  soluble  in  water.  A  salt  much  utilized  in  analysis 
is  ammonium-magnesium  phosphate,  Mg  (NH4)PO4.  This 
is  difficultly  soluble  in  water,  and  may  therefore  be  used 
either  for  the  purpose  of  detecting  magnesium  or  phos- 
phoric acid.  In  order  to  produce  this  salt,  ammonia 
and  some  ammonium  salt,  together  with  a  soluble  mag- 
nesium salt,  must  be  added  to  a  soluble  phosphate. 
If  ammonia  alone  were  added  to  a  solution  containing  a 
magnesium  salt,  magnesium  hydroxide  would  be  precipi- 
tated : 

MgS04  +  2NH4OH  =  Mg(OH)2  +  (NH4)2SO4. 


BORATES—  SILICATES—  MAGNESIUM  SILICIDE.     561 

'•   •' 

With  ammonium  salts,  however,  magnesium  salts  form 
compounds,  which  are  not  decomposed  on  the  addition 
of  ammonia.  When  a  soluble  phosphate  is  added,  the 
difficultly  soluble  ammonium-magnesium  salt  is  thrown 
down.  When  heated,  this  salt  loses  ammonia,  then 
water,  and  is  converted  into  magnesium  pyrophosphate  : 

Mg(NH4)P04  =  MgHP04  +  NH3  ; 
SMgHPO,  =  Mg,PA  +  H,0. 


The  corresponding  salt  of  arsenic  acid,  Mg(NH4)AsO4,  is 
very  similar  to  the  phosphate,  and  on  account  of  its  in- 
solubility it  is  also  used  in  chemical  analysis. 

Borates.  —  A  borate  of  magnesium  together  with  mag- 
nesium chloride  occurs  in  nature,  and  is  known  as  bora- 
cite.  It  has  the  composition  expressed  by  the  formula 
2Mg3B8O15  +  MgCl2.  The  borate,  Mg3B8O15,  is  derived 
from  the  acid,  H6B8O15,  which  is  related  to  normal  boric 
acid,  as  is  shown  by  the  equation 


Silicates.  —  The  simplest  silicate  of  magnesium  found 
in  nature  is  olivine,  which  is  represented  by  the  formula 
Mg2SiO4.  It  is  the  neutral  salt  of  normal  silicic  acid. 

Serpentine  is  derived  from  the  acid,  O<«i(Qjj\3>  and  has 

the  composition  Mg3Si2O7  +  2H2O. 

Magnesium  Silicide,  Mg2Si,  is  made  by  heating  together 
magnesium  chloride,  sodium  fluosilicate,  sodium  chlo- 
ride, and  sodium.  Under  these  circumstances  the  so- 
dium sets  magnesium  free  from  the  chloride,  and  silicon 
from  the  fluosilicate.  Both  unite  to  form  magnesium 
silicide.  When  treated  with  hydrochloric  acid  it  gives 
silicon  hydride,  SiH4,  and  hydrogen  : 

MgaSi  +  4HC1  =  2MgCl2  +  SiH4. 

The  liberation  of  hydrogen  is  due  to  the  presence  of 
an  excess  of  magnesium. 

Reactions   of  Magnesium   Salts  which   are   of  Special 


562  INORGANIC  CHEMISTRY. 

• 

Value  in  Chemical  Analysis. — Soluble  hydroxides  (KOH, 
NaOH,  NH4OH)  precipitate  magnesium  hydroxide.  If 
ammonium  chloride  is  present  ammonia  does  not  pre- 
cipitate the  hydroxide. 

Di-sodium  phosphate  with  ammonia  and  ammonium 
chloride  precipitates  ammonium-magnesium  phosphate 
from  the  solution  of  a  magnesium  salt. 

Sodium  and  potassium  carbonates  precipitate  basic  mag- 
nesium carbonate. 

ERBIUM,  E  (At.  Wt.  165.06). 

General. — As  regards  the  position  of  erbium  in  the  pe- 
riodic system,  a  final  statement  cannot  as  yet  be  made. 
According  to  its  atomic  weight,  assuming  it  to  be  165.06,  it 
falls  in  the  second  family.  On  the  other  hand,  the  com- 
position of  its  compounds  seems  to  indicate  rather  that  it 
belongs  in  the  third  family,  as  it  resembles  aluminium 
in  some  respects.  It  occurs  in  some  rare  minerals,  as 
cerite,  gadolinite,  euxenite,  and  orthite,  which  are  found 
in  Sweden  and  Greenland.  It  is  always  accompanied 
by  other  rare  metals,  a  few  of  which  have  been  studied 
with  care.  Among  these  may  be  mentioned  lanthanum, 
cerium,  didymium,  and  scandium.  These  metals  will 
be  treated  of  in  later  chapters.  It  need  only  be  said 
further  in  regard  to  erbium,  that  our  knowledge  con- 
cerning it  is  as  yet  quite  imperfect,  and  the  cause  of 
this  is  to  be  found  in  the  fact  that  the  minerals  in  which 
it  occurs  are  exceedingly  complex,  and  it  is  therefore 
very  difficult  to  separate  the  various  metals  present.  It 
appears  that  the  formula  of  the  oxide  of  erbium  is  E2O3. 
If  this  is  so,  it  is  in  this  respect  like  aluminium  oxide, 

AlA. 

It  appears  probable  that  the  oxide  of  erbium  that  is 
generally  obtained  from  the  complex  minerals  named  is 
a  mixture  of  oxides  of  rare  elements.  There  is  good  evi- 
dence of  the  presence  of  thulium  and  of  holmium,  though 
it  is  still  an  open  question  whether  these  as  well  as  erbium 
may  not  be  capable  of  further  decomposition. 


CHAPTER  XXVII. 

ELEMENTS  OF  FAMILY  III.  GROUP  A  : 
ALUMINIUM— SCANDIUM— YTTRIUM— YTTERBIUM- 
SAMARIUM— HELIUM. 

General. — There  is  in  some  respects  a  resemblance 
between  boron  and  the  principal  member  of  this  group  ; 
but  as  boron  acts  almost  exclusively  as  an  acid-forming 
element,  it  was  taken  up  in  connection  with  the  elements 
of  Family  V,  Group  B,  or  the  nitrogen  group.  Atten- 
tion was,  however,  called  to  the  fact  that  the  analogy 
between  these  elements  and  boron  is  but  slight.  The 
points  of  resemblance  between  boron  and  the  members 
of  Family  III,  Group  A,  will  be  pointed  out  below. 
The  principal  member  of  this  group  is  aluminium.  The 
others  are  all  rare,  and  some  have  been  but  imperfectly 
studied,  owing  to  serious  difficulties  in  the  way  of  ob- 
taining their  compounds  in  pure  condition.  They  are 
trivalent  in  their  compounds,  the  general  formulas  being 
such  as  the  following  : 

MCI,,  M(OH)a,  M(N03)3,  M,(S04),,  M/CO,)*  MPO4)  etc. 

Aluminium  oxide  is  weakly  basic,  and  somewhat  acidic, 
though  less  so  than  boron.  Aluminium  hydroxide  has 
the  power  to  neutralize  most  acids,  and  also  to  form  salts 
with  strong  bases.  Boron  oxide,  on  the  other  hand,  has 
scarcely  any  basic  properties,  though  it  does  form  a  few 
extremely  stable  compounds,  in  which  the  boron  replaces 
the  hydrogen  of  acids.  (See  Boron  Phosphate,  p.  356.) 

ALUMINIUM,  Al  (At.  Wt.  26.91). 

Occurrence. — Aluminium  is  an  extremely  important 
element  in  nature  and  in  the  arts.  It  occurs  very 

(563) 


564  INORGANIC  CHEMISTRY. 

widely  distributed,  and  very  abundantly  in  many  different 
forms  of  combination.  Among  them  are  feldspar,  mica, 
cryolite,  bauxite.  Feldspar  is  a  silicate  of  aluminium 
and  potassium  of  the  formula  AlKSi3O8.  Mica  is  a  gen- 
eral name  applied  to  a  large  number  of  minerals  which 
are  silicates  of  aluminium  and  some  other  metal,  as  po- 
tassium, lithium,  magnesium,  etc.  The  simplest  form 
of  mica  is  that  represented  by  the  general  formula 
KAlSiO4,  according  to  which  the  mineral  is  a  salt  of 
orthosilicic  acid,  Si(OH)4.  Cryolite  is  a  double  fluoride 
of  aluminium  and  sodium,  or  the  sodium  salt  of  fluo- 
aluminic  acid,  Na3AlF6.  Bauxite  is  a  hydroxide  of 
aluminium  in  combination  with  a  hydroxide  of  iron. 
Besides  in  the  above  forms,  aluminium  occurs  in  the 
products  of  decomposition  of  minerals.  One  of  the  most 
important  of  these  is  clay,  which  is  found  in  all  condi- 
tions of  purity — from  the  white  kaoline  to  ordinary 
dark-colored  clay.  Kaoline  is  the  aluminium  salt  of 
orthosilicic  acid  of  the  formula  Al4(SiO4)3  -f-  4H2O.  Alu- 
minium silicate  is  found  in  all  soils,  but  is  not  taken  up 
by  plants,  and  does  not  find  entrance  into  the  animal 
body.  The  name  aluminium  has  its  origin  in  the  fact 
that  the  salt  alum  was  known  at  an  early  date,  and  the 
metal  was  afterwards  isolated  from  it. 

Preparation. — The  preparation  of  aluminium  on  the 
large  scale  presents  a  problem  of  the  highest  importance 
to  the  human  race.  The  element  has  properties  which 
adapt  it  to  many  uses  to  which  iron  is  put,  and  for  many 
purposes  it  has  many  advantages  over  iron.  Further,  we 
are  supplied  by  nature  with  unlimited  quantities  of  the 
compounds  of  aluminium,  which  are  distributed  every- 
where over  the  earth.  While,  however,  iron,  lead,  tin,, 
copper,  and  other  metals  can  be  isolated  from  their 
natural  compounds  without  serious  difficulty,  aluminium, 
which  is  more  abundant  than  any  of  them,  and  in  many 
respects  more  valuable  than  any  of  them,  is  locked  in 
its  compounds  so  firmly,  that  it  is  only  by  comparatively 
complicated  and  expensive  methods  that  it  can  be  iso- 
lated ;  and  up  to  the  present  it  cannot  be  made  at  a 
price  sufficiently  low  to  bring  it  into  common  use.  At 


ALUMINIUM.  565 

the  same  time  work  is  constantly  in  progress  with  refer- 
ence to  this  important  practical  problem,  and  it  seems 
probable  that  through  a  thorough  study  of  the  Jaws  of 
chemistry  some  method  for  the  cheap  preparation  of  alu- 
minium on  the  large  scale  will  eventually  be  discovered. 
The  first  method  devised  for  the  preparation  of  alu- 
minium on  the  large  scale  consisted  in  heating  aluminium 
chloride  with  sodium.  The  chloride  was  heated  to  boil- 
ing in  a  retort ;  the  vapor  passed  through  a  vessel  contain- 
ing pieces  of  iron  heated  to  redness,  and  then  into  a  long 
tube  containing  sodium.  Instead  of  aluminium  chloride, 
the  double  chloride  of  aluminium  and  sodium,  which  is 
more  easily  prepared  in  the  dry  condition,  is  now  used. 
The  double  chloride  and  cryolite  are  heated  together 
with  sodium  in  a  properly  constructed  furnace.  It  is, 
further,  possible  to  prepare  aluminium  by  electrolysis 
of  the  chloride  or  of  the  double  chloride  above  men- 
tioned ;  and  the  oxide  can  be  reduced  by  mixing  it  with 
charcoal  and  passing  the  current  from  a  powerful  dy- 
namo-machine through  it.  By  the  latter  method  an 
alloy  of  aluminium  and  copper  is  now  prepared,  but  the 
preparation  of  aluminium  alone  by  this  method  does  not 
appear  to  be  entirely  successful.  New  methods  for  the 
preparation  of  the  metal  are  constantly  being  devised, 
and  the  price  is  constantly  being  lowered.  The  latest 
method  of  promise  consists  in  the  electrolysis  of  alu- 
minium oxide,  in  the  form  of  corundum,  in  a  bath  of 
molten  cryolite  contained  in  a  carbon  crucible.  A  large 
number  of  patents  have  been  issued,  covering  methods 
for  the  preparation  of  aluminium;  but  these  are  fre- 
quently so  imperfectly  described,  and  the  evidence  of 
their  value  so  unsatisfactory,  that  it  is  difficult  to  pass 
judgment  upon  them.  Until  recently  the  commercial 
preparation  of  aluminium  has  appeared  to  be  intimately 
connected  with  that  of  the  commercial  preparation  of 
sodium  ;  but,  in  view  of  the  success  of  the  electrolytic 
method,  this  is  no  longer  the  case. 

Properties. — The  color  of  aluminium  is  like  that  of 
tin,  and  it  has  a  high  lustre.  It  is  very  strong,  .and  yet 
malleable.  It  is  lighter  than  most  metals  in  common  use, 


566  INORGANIC  CHEMISTRY. 

its  specific  gravity  being  2.5  to  2.7  according  to  the  con- 
dition, while  that  of  iron  is  7.8,  that  of  silver  10.57,  that 
of  tin  7,3,  and  that  of  lead  11.37.  It  does  not  change  in 
dry  or  in  moist  air  ;  and  in  the  compact  form  it  does  not 
act  upon  water  even  at  elevated  temperatures.  It  melts 
at  about  700°,  which  is  higher  than  the  melting-point  of 
zinc,  and  lower  than  that  of  silver.  Hydrochloric  acid 
dissolves  it  with  ease,  forming  aluminium  chloride.  At 
the  ordinary  temperatures  nitric  and  sulphuric  acids  do 
not  act  upon  it ;  at  higher  temperatures,  however,  action 
takes  place,  and  the  corresponding  salts  are  formed.  It 
dissolves  in  solutions  of  the  caustic  alkalies,  forming  the 
so-called  aluminates.  It  reduces  many  oxides  when 
heated  with  them  to  a  sufficiently  high  temperature ; 
and  is  used  in  the  preparation  of  boron  and  silicon. 

Applications. — The  metal  is  used  to  a  considerable 
extent  in  the  preparation  of  ornaments,  and  of  useful 
articles  in  which  lightness  is  a  matter  of  importance,  as 
in  telescopes  and  opera-glasses.  An  alloy  with  a  small 
percentage  of  silver  is  used  for  the  beams  of  chemical 
balances.  Aluminium  bronze,  which  is  an  alloy  with 
copper,  is  also  used  quite  extensively.  It  will  be  again 
referred  to  under  Copper. 

Aluminium  Chloride,  A1C13. — When  aluminium  hydrox- 
ide is  dissolved  in  hydrochloric  acid  a  solution  of  alu- 
minium chloride  is  formed,  and  from  this  solution  a 
compound  of  the  formula  A1C13  +  6H2O  can  be  obtained 
in  crystallized  form.  Like  calcium  and  magnesium  chlo- 
rides, this  salt  is  deliquescent.  When  heated  to  drive  off 
the  water  the  salt  conducts  itself  like  magnesium  chlo- 
ride, but  the  decomposition  into  the  oxide  and  hydro- 
chloric acid  takes  place  more  easily  than  that  of  mag- 
nesium chloride.  The  reaction  is  represented  by  the 
equation 

2A1C13  +  3H20  =  A1303  +  6HC1. 

The  dry  chloride  is  prepared  by  the  same  method  as 
that  used  in  the  preparation  of  silicon  chloride  and  boron 
chloride,  viz.,  by  passing  chlorine  over  a  heated  mixture 
of  the  oxide  and  carbon.  The  chloride,  being  volatile, 


ALUMINIUM  CHLORIDE.  567 

sublimes,  and  is  deposited  in  the  cool  part  of  the  vessel, 
when  pure,  as  a  white  laminated  crystalline  mass.  Gen- 
erally, however,  it  is  more  or  less  colored  in  consequence 
of  the  presence  of  impurities.  When  exposed  to  the  air 
it  attracts  moisture  and  gives  off  hydrochloric  acid.  It 
dissolves  in  water  very  easily,  with  a  marked  evolution  of 
heat,  but,  from  what  was  said  above,  it  is  evident  that  it 
cannot  be  obtained  from  this  solution  again  by  evapora- 
tion. It  is  volatile  without  change.  The  specific  gravity 
of  its  vapor  has  been  determined  by  different  observers, 
and,  unfortunately,  with  different  results.  According  to 
Deville  and  Troost,  it  is  such  as  to  lead  to  the  formula 
A12C16.  Quite  recently,  however,  Nilson  and  Pettersson 
have  found  it  to  correspond  to  that  required  by  the  for- 
mula A1C13,  their  determinations  having  been  made  at  a 
higher  temperature  than  those  of  Deville  and  Troost. 
Still  later  determinations  by  Crafts  again  lead  to  the 
formula  A12C16.  Upon  the  basis  of  the  determinations  by 
Deville  and  Troost,  chemists  have  for  some  time  past  used 
the  formula  AlaCl6  to  represent  the  compound.  Accord- 
ing to  this,  aluminium  would  appear  to  be  quadrivalent, 
as  represented  in  the  following  formula  for  the  chloride  : 
Ck  /Cl 

C1-)A1-A1(—C1 .  On  the  other  hand,  in  a  compound 
CK  \C1 

made  by  replacing  the  chlorine  of  this  chloride  by  certain 
organic  groups  the  aluminium  appears  to  be  trivalent,  as 
represented  in  the  formula  A1(CH3)3,  in  which  the  group 
CH3,  known  as  methyl,  is  univalent.  Further,  the  posi- 
tion of  aluminium  in  the  periodic  system  makes  it  appear 
extremely  probable  that  it  is  trivalent,  and  not  quadriv- 
alent. What,  then,  is  the  explanation  of  the  discrepancy 
above  noted  in  the  evidence  regarding  the  constitution 
of  the  chloride  ?  When  we  come  to  examine  the  conduct 
of  aluminium  chloride  towards  the  chlorides  of  other 
metals,  and  find  with  what  ease  it  forms  double  chlorides, 
it  seems  not  improbable  that  aluminium  chloride  itsolf, 
at  ordinary  temperatures,  and  even  in  the  form  of  vapor 
at  lower  temperatures,  may  be  a  compound  of  the  same 
order  as  the  double  chlorides.  It  has  been  suggested 


568  INORGANIC  CHEMISTRY. 

that  in  these  compounds  chlorine  is  probably  in  com- 
bination with  chlorine,  as  fluorine  is  with  fluorine  in  hy- 
drofluoric acid,  in  such  a  way  that  two  chlorine  atoms 
can  exert  a  linking  function  between  two  other  atoms. 

/Cl 
Just  as  there  is  a  compound  of  the  formula  A1—-C1 


so  it  is  possible  that  aluminium  chloride  may  have  the 
constitution  represented  by  the  formula  A1^(C12)^)A1, 

in  which  the  aluminium  is  trivalent.  By  replacing  the 
chlorine  in  a  compound  of  this  constitution  by  groups 
like  methyl,  which  cannot  exert  the  linking  function, 
the  product  would  not  be  a  double  compound.  Further, 
by  heating  a  compound  of  this  constitution  it  would 
probably  dissociate  into  two  molecules  of  the  simple 
compound  A1C13,  and  it  would  be  this  which  comes  into 
play  in  chemical  reactions.  In  view  of  the  conflicting 
state  of  the  evidence  and  the  plausibility  of  the  above 
explanation,  the  formula  for  aluminium  chloride  used 
here  is  the  simpler  one.  By  means  of  it  and  similar  for- 
mulas for  the  other  compounds  of  aluminium,  the  reac- 
tions of  the  element  can  be  expressed  somewhat  more 
easily  and  probably  just  as  truthfully  as  by  means  of  the 
more  complicated  formula. 

Chloroaluminates  or  Double  Chlorides  of  Aluminium 
and  Analogous  Compounds. — These  compounds  have 
been  repeatedly  referred  to,  and  but  very  little  need  be 
added  to  what  has  already  been  said  concerning  them. 
In  general,  the  chloride,  bromide,  and  iodide  of  aluminium 
combine  with  the  chlorides,  bromides,  and  iodides  of  the 
most  strongly  marked  metals,  such  as  potassium  and 
sodium.  Those  with  potassium  and  sodium  have  the  for- 

XSL 
mulas  A1C13.KC1  and  AlCl3.NaCl,  or  probably  Al(-Cl 

\C1,)K 
/Cl 
and  A1^-C1         .      The  fluoride  forms   two   compounds 

\Cla)Na 
with  potassium  fluoride  and  two  with  sodium  fluoride. 


ALUMINIUM  HYDROXIDE.  569 

These  have  the  composition    represented   by  the    for- 
mulas A1F3.2KF,  AlF3.2NaF,  and  A1F3.3KF,  AlF3.3NaF, 

XF 
and    the   constitution    expressed  thus,   A1(-(F2)K    and 


A1(~(F,)K  .     The   tri-sodium  fluoaluminate  is  the  min- 


eral  cryolite,  which  occurs  in  such  large  quantity  as  to 
be  exported,  and  form  the  starting-point  in  the  prepa- 
ration of  aluminium  and  even  sodium  compounds.  A 
method  for  making  sodium  carbonate  from  cryolite  has 
already  been  described.  Its  use  in  the  preparation  of 
aluminium  compounds  will  be  taken  up  as  far  as  may  be 
necessary  in  this  chapter. 

Besides  the  compounds  with  metallic  chlorides,  alu- 
minium chloride  also  forms  compounds  with  the  chlorides 
of  the  acid-forming  elements.  Such,  for  example,  are 
the  compounds  with  sulphur  tetrachloride  and  with 
phosphorus  pentachloride.  These  have  the  composition 
represented  by  the  formulas  (A1C13)2SC14  and  A1C13.PC15. 
The  latter  may  be  the  chlorine  analogue  of  aluminium 
phosphate,  A1PO4.  If  the  oxygen  in  the  phosphate 
should  be  replaced  by  an  equivalent  quantity  of  chlorine 
the  result  would  be  a  compound  of  the  formula  A1PC18, 
which  is  that  of  the  above  compound.  These  double 
chlorides,  like  the  chlorides  of  the  acid-forming  elements 
in  general,  are  easily  decomposed  by  water,  yielding  the 
corresponding  oxygen  compounds.  A  compound  inter- 
mediate between  the  oxygen  and  the  chlorine  compounds 
is  that  formed  by  the  combination  of  aluminium  chloride 
with  phosphorus  oxychloride,  which  is  represented  by 
the  formula  A1POC16,  or  A1C13.POC13.  This  may  be  re- 
garded as  aluminium  phosphate,  in  which  three  of  the 
oxygen  atoms  have  been  replaced  by  six  chlorine  atoms. 

Aluminium  Hydroxide,  A1(OH)3.  —  Normal  aluminium 
hydroxide,  A1(OH)3,  occurs  in  nature  as  the  mineral  hy- 
drargillite.  It  is  precipitated  from  a  solution  of  alu- 
minium chloride  by  ammonia  : 

A1C1,  +  3NH4OH  =  A1(OH)3  +  3NH4C1. 


570  INORGANIC  CHEMISTRY. 

Obtained  by  precipitation  it  forms  a  gelatinous  mass, 
which  is  suggestive  of  starch-paste,  and  it  is  on  this 
account  extremely  difficult  to  wash  it  completely  free 
from  the  substances  in  the  solution.  It  dries  in  the  air, 
forming  a  gummy  substance  which  has  the  composition 
A1(OH)8.  When  heated  under  proper  conditions  it  loses 
water,  and  forms  the  compound  A1O2H  : 

A1(OH)3  =  A1O.OH  +  H2O. 

This  compound  is  found  in  nature  as  the  mineral  dias- 
pore.  If  heated  to  a  higher  temperature  it  is  converted 
into  the  oxide,  A12O3  : 

2A1(OH)3  =  A1203  +  3H20. 

In  the  conduct  of  the  chloride  and  of  the  hydroxide 
aluminium  exhibits  a  certain  resemblance  to  boron.  The 
acidic  character  of  the  latter  is,  however,  more  strongly 
marked  than  that  of  the  former.  Boron  chloride  is  more 
easily  decomposed  by  water  than  aluminium  chloride, 
and,  as  the  decomposition  takes  place  at  the  ordinary 
temperature,  the  product  is  the  hydroxide  instead  of  the 
oxide,  as  in  the  case  of  aluminium.  The  hydroxide, 
B(OH)3,  readily  loses  water  and  forms  metaboric  acid, 
which  in  composition  is  analogous  to  diaspore  ;  and  at  a 
higher  temperature  the  oxide,  B2O3,  is  formed. 

Besides  the  normal  hydroxide,  A1(OH)3,  and  that  of 
the  formula  AIO(OH),  there  is  a  third  one  known.  This 
has  the  composition  A12O(OH)4,  and,  as  is  plain,  is  derived 
from  two  molecules  of  the  normal  hydroxide  by  loss  of 
one  molecule  of  water  : 

/OH  TT/^  r\TT 

H6>A1-°-A1<OH  +  H>°- 


This  has  been  obtained  in  solution  ;  or,  rather,  it  has 
been  obtained  by  evaporation  of  a  solution  of  hydroxide 
made  by  continued  boiling  of  a  solution  of  basic  acetate 
of  aluminium  which  decomposes  into  hydroxide  and 
acetic  acid,  the  latter  then  evaporating.  From  this 
solution,  by  evaporation  in  a  water-bath,  the  above  hy- 


ALUMINATES.  571 

droxide  is  obtained.  As  already  stated,  bauxite  is,  in  all 
probability,  a  compound  of  this  constitution  combined 
with  a  similar  hydroxide  of  iron.  A  hydroxide  of  the 
same  composition  is  obtained  when  a  solution  of  the 
normal  hydroxide  in  caustic  soda  is  boiled  with  am- 
monium chloride.  The  precipitate  formed  in  this  way  is 
not  gelatinous,  and,  when  dried,  it  has  the  composition 
A120(OH), 

The  preparation  of  aluminium  hydroxide  from  natural 
compounds  of  the  element  is  based  upon  the  fact  that 
aluminium  oxide  forms  with  sodium  a  soluble  compound, 
and  that  this  is  decomposed  by  carbon  dioxide  with  pre- 
cipitation of  the  hydroxide.  The  sodium  compound 
formed  has  probably  the  composition  Al(ONa)8,  being  a 
salt  of  the  normal  hydroxide.  When  this  is  treated  in 
solution  with  carbon  dioxide,  the  decomposition  takes 
place  as  represented  in  this  equation : 

2Al(ONa)3  +  3CO2  +  3H2O  =  3Na2CO3  +  2A1(OH)3. 

When  cryolite  is  ignited  with  lime,  the  products  are 
probably  calcium  fluoride,  sodium  oxide,  and  another 
variety  of  sodium  aluminate  : 

Na8AlF6  +  3CaO  =  NaA102  +  Na2O  +  3CaF,. 

When  the  mass  is  treated  with  water,  the  calcium  fluor- 
ide remains  undissolved,  while  the  sodium  and  aluminium 
form  the  compound  Al(ONa)3.  This  undergoes  decompo- 
sition, as  above  represented,  when  treated  with  carbon 
dioxide.  Two  valuable  products — aluminium  hydroxide 
and  sodium  carbonate — are  thus  obtained. 

In  order  to  prepare  the  hydroxide  from  bauxite,  this 
is  heated  to  a  high  temperature  with  sodium  carbonate. 
Water  extracts  sodium  aluminate,  from  which  the  hy- 
droxide is  precipitated  by  means  of  carbon  dioxide. 

Aluminium  hydroxide  forms  the  material. for  the  prep- 
aration of  aluminium  salts ;  as,  the  chloride,  sulphate, 
alum,  etc. 

Aluminates. — When  sodium  or  potassium  hydroxide  is 
added  to  a  solution  of  an  aluminium  salt,  aluminium  hy- 
droxide is  at  first  precipitated,  but  an  excess  of  the  re- 


572  INORGANIC  CHEMISTRY. 

agent  used  dissolves  the  precipitate.  This  action  is  the 
same  in  character  as  that  which  takes  place  in  the  case 
of  glucinum,  and  is  due  to  the  acidic  character  of  alu- 
minium hydroxide.  It  is  probable  that  in  solution  the 
action  with  potassium  and  sodium  hydroxides  is  of  the 
same  kind  as  represented  in  the  equations 

A1(OH)3  +  3KOH    =A1(OK)3    +  3H,O,  and 
A1(OH),  +  3NaOH  =  Al(ONa)3  +  3H2O. 

On  evaporating  the  solution  of  the  potassium  salt,  how- 
ever, the  product  obtained  has  the  formula  A1O.OK, 
and  is  plainly  the  salt  of  the  hydroxide  A1O.OH,  which 
may  be  called  meta-aluminic  acid,  to  suggest  its  analogy 
to  metaboric  acid,  BO. OH.  When  aluminium  hydroxide 
and  sodium  carbonate  are  melted  together,  the  salt 
AlO.ONais  formed,  as  has  been  shown  by  determining 
the  amount  of  carbon  dioxide  given  off  when  a  known 
weight  of  the  hydroxide  is  employed.  When,  however, 
the  solution  of  the  hydroxide  in  caustic  soda  is  evap- 
orated, the  salt  Al(ONa)3  is  deposited. 

These  salts  are  very  unstable,  though  their  solutions 
can  be  boiled  without  undergoing  decomposition.  Car- 
bon dioxide  decomposes  them  at  once  with  precipitation 
of  aluminium  hydroxide,  as  was  stated  in  describing  the 
method  for  the  preparation  of  the  hydroxide  from  cryolite 
and  from  bauxite.  Similar  salts  are  formed  with  calcium 
and  barium.  Among  them  may  be  mentioned  those  of  the 
following  formulas :  Ca3(AlO3)2,  Ca(AlO2)2,  Ba3(AlO3)2,  and 
Ba(AlO2)2.  The  calcium  salts  are  insoluble  in  water,  and 
some  of  them  become  hard  in  contact  with  water.  They 
are  therefore  of  importance  in  the  manufacture  of  hy- 
draulic cements.  The  barium  salts  are  soluble  in  water. 

Many  aluminates  occur  in  nature,  forming  the  import- 
ant group  of  minerals  known  as  the  spinels.  Of  these,  spi- 
nel itself  is  the  magnesium  salt  of  the  hydroxide  A1O.OH, 

and    is    represented    by    the    formula   ^Q *Q>Mg,  or 

Mg(AlO,)2.  Chrysoberyl  is  the  corresponding  glucinum 
salt  G1(A1O2)2;  and  gahnite  is  the  zinc  salt  Zn(AlO2)2. 
These  salts  are  extremely  stable,  differing  markedly  in  this 


ALUMINATES.  573 

respect  from  those  above  referred  to,  which  are  made  in 
the  laboratory.  They  are  decomposed  by  heating  them, 
in  finely  powdered  condition,  with  primary  or  acid  po- 
tassium sulphate,  the  action  of  which  was  described  on 
p.  499.  As  will  be  seen  farther  on,  there  are  other  salts 
similar  to  the  aluminates  in  structure  which  occur  in 
nature.  Among  these  there  may  be  mentioned  here 
chromic  iron,  or  chromite,  which  is  an  iron  salt  of  a  hy- 
droxide of  chromium  of  the  formula  CrO.OH.  The  salt 
is  to  be  regarded  as  made  up  according  to  the  formula 

CrO  O>Fe>  or  Fe(Cr°03-  Further,  magnetic  oxide  of 
iron  or  magnetite,  Fe3O4,  is  regarded  as  belonging  to 
the  same  group,  and  its  constitution  represented  thus  : 

,  or  Fe(FeO2)2;  and  there  is  also  a  compound 


of  magnesium,  ^  r\  'o>Mg.     For  the  sake  of  empha- 


sizing these  analogies,  the  formulas  of  the  compounds 
above  mentioned  are  here  presented  in  tabular  form  : 

Potassium  aluminate,     .     .     A1O.OK 
Sodium  aluminate,     .     .     .     AlO.ONa 

Calcium  aluminate,   .     .     .     ~\]r\  *Q>Ca 
Barium  aluminate,     .     .     .  ' 


Spinel,     .     .    .....     .     .     . 

Chrysoberyll,   .....     AlO.'o>G1 


Gahnite,  ....... 

Chromite,     ......  ' 


Magnetite,    ......     FeO.'o>Fe 

Magnesio-ferrite,  .... 

There   is   a   highly  instructive   analogy  between  the 
aluminates  and  the  double  chlorides  and  other  similar 


574  INORGANIC   CHEMISTRY. 

compounds.  In  general,  aluminium  hydroxide  acts  upon 
the  hydroxides  of  the  strongest  base-forming  elements 
to  form  aluminates.  So,  also,  aluminium  chloride  acts 
upon  the  chlorides  of  the  strongest  base-forming  ele- 
ments to  form  double  chlorides.  By  melting  together 
aluminium  hydroxide  and  potassium  or  sodium  hydrox- 
ide, compounds  of  the  formulas  A1O.OK  and  AlO.ONa 
are  formed.  So,  also,  by  melting  together  aluminium 
chloride  and  potassium  or  sodium  chloride,  compounds 
of  the  formulas  A1C14K  and  AlCl4Na  are  formed.  Com- 
paring these  oxygen  and  chlorine  compounds,  it  is  clear 
that  they  are  analogous.  If  the  oxygen  of  the  former 
is  replaced  by  an  equivalent  quantity  of  chlorine,  the 
chlorine  compounds  result  : 

EA1O2  KA1C14 

NaAlO,  NaAlCl4 

Or,  if  their  constitutional  formulas  are  written  in  accord- 
ance with  the  views  already  expressed  regarding  the 
double  chlorides,  the  analogy  is  also  seen,  thus  : 

O  /Cl 

AlCl 


\C12)K 

A1\ONa  A1^C1 

\Cla)Na. 

The  compounds  of  the  same  order  as  cryolite  have  their 
analogues  in  such  oxygen  compounds  as  Al(ONa)8,  etc., 
as  is  shown  by  the  following  formulas  : 

Na3AlO8  ;  Na3AlF6  ; 

ONa 


f  -(F2)Na  . 


It  is  not  improbable  that  by  fusion  with  other  chlorides 
besides  those  of  potassium  and  sodium,  aluminium  chlo- 
ride will  be  found  to  yield  other  double  chlorides  analo- 
gous to  the  spinels.  According  to  what  was  said  in 


ALUMINIUM  OXIDE.  5~5 

discussing  the  subject  of  double  chlorides  in  general, 
three  series  of  such  salts  may  be  looked  for,  correspond- 
ing to  the  formulas 

/Cl  /Cl  /(C12)M    - 

AlfCl        ,  A1(--(C12)M,       and        Alf  (C12)M ; 

\C12)M  \C12)M  \C12)M 

and  representatives  of  all  these  classes  are  known.  Oxy- 
gen compounds  corresponding  to  the  first  and  last  of 
these  have  been  mentioned.  As  an  example  of  an  oxygen 
compound  corresponding  to  the  second  one,  barium 
aluminate,  of  the  formula  Ba2Al2O5,  may  be  cited. 

Aluminium  Oxide,  A12O3. — As  has  been  stated,  the 
oxide  is  formed  by  heating  the  hydroxide.  It  is  found 
in  nature  in  the  form  of  ruby,  sapphire,  and  corundum. 
The  natural  variety  is  extremely  hard ;  and  granular 
corundum,  which  is  known  as  emery,  is  used  for  polish- 
ing. The  red  color  of  the  ruby  is  caused  by  the  presence 
of  a  trace  of  a  chromium  compound  ;  while  the  blue  color 
of  the  sapphire  is  probably  due  to  the  presence  of 
a  trace  of  a  cobalt  compound.  Aluminium  oxide  is 
infusible  in  the  hottest  furnace  fire,  but  it  melts  in  the 
flame  of  the  oxyhydrogen  blow-pipe,  and  on  cooling  it 
becomes  crystalline.  By  mixing  it  with  various  easily 
fusible  substances  and  heating,  it  is  obtained  in  the  form 
of  crystals,  and  by  adding  certain  metallic  oxides  these 
crystals  can  be  colored.  In  this  way  artificial  rubies  and 
sapphires  have  been  prepared,  which  have  all  the  prop- 
erties of  the  natural  ones.  When  the  oxide  is  moistened 
with  a  few  drops  of  a  solution  of  cobaltous  nitrate  and 
then  ignited,  it  turns  blue.  This  fact  is  taken  advantage 
of  in  chemical  analysis  for  the  purpose  of  detecting  alu- 
minium. When  the  oxide  is  made  by  gently  igniting 
the  hydroxide,  it  dissolves  in  strong  acids.  If,  however, 
it  is  heated  to  a  high  temperature,  acids  will  not  dissolve 
it.  The  natural  varieties  of  the  oxide,  further,  are  not 
soluble  in  acids.  By  fusion  with  acid  potassium  sulphate 
insoluble  aluminium  oxide  is  converted  into  a  soluble 
compound. 


576  INORGANIC  CHEMISTRY. 

Aluminium  Sulphate,  A12(SO4)3. — This  salt  is  made  by 
dissolving  the  hydroxide  of  aluminium  in  dilute  sulphuric 
acid,  and  evaporating  to  crystallization,  when  a  salt  of 
the  composition  A12(SO4)3  +  18H2O  is  deposited.  When 
heated  the  salt  loses  its  water  of  crystallization,  and,  if 
the  temperature  is  raised  to  that  of  red  heat,  the  anhy- 
drous salt  is  decomposed  with  loss  of  sulphur  trioxide 
and  formation  of  aluminium  oxide.  This  decomposition 
is,  however,  not  complete.  The  sulphate  is  manufactured 
on  the  large  scale  for  various  purposes,  as,  for  example, 
for  a  mordant,  for  sizing  paper,  etc. 

Basic  Aluminium  Sulphates. — A  solution  of  ordinary 
aluminium  sulphate  has  an  acid  reaction,  and  has  the 
power  to  dissolve  metals,  such  as  zinc,  and  hydroxides, 
such  as  aluminium  hydroxide.  When  a  solution  of  the 
sulphate  is  treated  with  the  hydroxide,  a  basic  salt  of 
the  formula  A12O(SO4)2  +  H2O  is  formed.  This  should 

(o  so 

probably  be  represented  by  the  formula  Al  •<  O          2  or 

(  OH 

A1(OH)SO4.  Another  basic  salt  has  the  formula 
(A1O)2SO4,  the  salt  being  derived  from  the  hydroxide, 

A1O.OH,  as  represented  thus :  ^JO  Q> SO2.     The  former 

salt  is  soluble  in  water.  When,  therefore,  a  solution  of 
sodium  or  ammonium  carbonate  is  added  to  a  solution 
of  ordinary  aluminium  sulphate,  the  first  portions  of  hy- 
droxide which  are  precipitated  redissolve  in  the  excess 
of  the  ordinary  salt.  There  are  other  basic  salts,  some 
of  which  occur  in  nature. 

Alums. — When  a  solution  of  aluminium  sulphate  is 
brought  together  with  a  solution  of  potassium  sulphate 
in  the  proportion  of  their  molecular  weights,  a  salt 
crystallizes  out  which  has  the  composition  represented 
by  the  formula 

KA1(S04)2  +  12H2O   or  K2S04  +  A13(SO4)3  +  24H,O. 

The  most  rational  view  which  has  been  expressed  re- 
garding this  compound  is  that  it  has  the  constitution 


ALUMS.  577 

,OK 
2<0 

O--A1,  with  perhaps  some  of  the  so-called  water 


of  crystallization  present  in  the  form  of  hydroxyl.  This 
salt,  which  has  long  been  known  under  the  name  of  alum, 
is  the  type  of  a  class  of  similar  compounds,  all  of  which 
are  called  alums.  These  alums  may  be  regarded  as 
derived  from  the  ordinary  form  by  replacing  the  potas- 
sium by  sodium,  ammonium,  or  any  other  member  of 
the  sodium  group,  besides  some  other  metals.  Thus  a 
series  of  alums  is  obtained,  of  which  the  following  are 
examples  : 


NaAl(SO4)s      H 
LiAl(S04)s 
(NH4)A1(S04)H 
CsAl(SO.X       H 
TIA1(SO.X 

-  12H2O  ; 
H12H20; 

i-  12H[o  ; 
1-  12H,0. 

Again,  alums  are  derived  from  the  ordinary  form  by  re- 
placing the  aluminium  by  some  other  elements  which 
have  the  power  to  form  compounds  resembling  those  of 
aluminium,  as,  for  example,  iron,  chromium,  and  man- 
ganese. Such  alums  are  those  represented  by  the  follow- 
ing formulas  : 

KFe(SO4)2  +  12H2O  ; 
KCr(S04)2  +12H20; 
KMn(SO4)2  +  12H2O. 

In  each  of  these,  again,  the  potassium  can  be  replaced  as 
in  the  case  of  ordinary  alum  ;  so  that  the  class  includes 
a  comparatively  large  number  of  salts.  All  have  certain 
properties  in  common.  They  are  all  soluble  in  water, 
and  all  crystallize  in  the  same  forms,  which  are  regular 
octahedrons  combined  with  cubes.  If  a  crystal  of  one 
alum  be  suspended  in  the  solution  of  any  other  one  it 
will  continue  to  grow.  They  are  all  strictly  isomorphous. 
The  principal  alums  containing  aluminium  are  those  of 


578  INORGANIC  CHEMISTRY. 

potassium  and  ammonium,  both  of  which  are  manufac- 
tured on  the  large  scale. 

Potassium  Alum,  Potassium  -  Aluminium  Sulphate, 
KA1(SO4)2  +  12HaO. — Ordinary  alum  is  found  in  nature 
in  some  volcanic  regions.  The  mineral  alunite,  which  is 
a  basic  salt  of  the  formula  .K(A1O)3(SO4)2  +  3H2O,  or 
perhaps  K[A1(OH)2]3(SO4)2,  occurs  in  larger  quantities. 
When  this  salt  is  heated  and  treated  with  water,  ordinary 
alum  dissolves,  and  is  easily  obtained  from  the  solution. 
Another  source  of  alum  is  alum  shale.  This  occurs  in 
large  quantities  in  nature,  and  consists  of  coal,  clay,  and 
iron  pyrites.  When  it  is  heated  in  contact  with  the  air 
the  coal  burns,  as  do  also  the  sulphur  and  pyrites,  and 
sulphuric  acid  is  formed.  When  allowed  to  lie  for  a  time 
in  contact  with  the  air  the  iron  pyrites  is  converted  into 
sulphate  and  sulphuric  acid.  The  latter  acts  upon  the 
clay  or  aluminium  silicate,  forming  aluminium  sulphate, 
from  which  alum  can  easily  be  made.  It  is  easier  to 
treat  the  shale  and  similar  substances  with  sulphuric 
acid  directly,  and  this  method  is  now  generally  employed. 

Alum  dissolves  readily  in  hot  water,  357.5  parts  of  the 
crystallized  salt  dissolving  in  100  parts  of  water  at  100°, 
At  0°  only  3.9  parts  dissolve,  and  at  the  ordinary  tem- 
perature about  12  parts.  It  crystallizes  beautifully  in 
regular  octahedrons,  occasionally  with  cube  faces  devel- 
oped on  them.  Under  some  circumstances  it  crystallizes 
in  cubes.  When  heated,  alum  melts  in  its  water  of 
crystallization,  and  if  heated  to  a  sufficiently  high  tem- 
perature the  water  passes  off,  leaving  burnt  alum.  Heated 
higher  the  salt  decomposes,  forming  aluminium  oxide 
and  potassium  sulphate,  and  finally  potassium  aluminate 
is  formed.  When  potassium  hydroxide,  ammonia,  or 
the  carbonate  of  potassium,  sodium,  or  ammonium,  is 
added  in  small  quantity  to  a  solution  of  alum  the  pre- 
cipitate first  formed  redissolves.  If  this  is  continued 
until  the  reaction  is  neutral,  or  until  a  point  is  reached 
beyond  which  the  addition  of  the  reagent  produces  a 
precipitate  which  does  not  redissolve,  there  is  then  con- 
tained in  the  solution  a  basic  compound,  known  as  basic 
alum,  which  probably  has  the  composition  K2(A12O)(SO4)3. 


ALUMINIUM  SILICATE.  579 

When  the  solution  is  boiled  the  salt  contained  in  it  is 
decomposed,  forming  ordinary  alum  and  another  basic 
alum  which  is  insoluble  : 

3[K,(AlaO)(SO.)3]  =  K(A10)s(SO,)a+  3KA1(SO.),  +  K,SO, 

The  insoluble  compound  is  known  as  insoluble  alum, 

Alum  crystallized  in  cubes  is  obtained  by  evaporating  a 
solution  to  which  some  sodium  or  potassium  carbonate 
has  been  added.  Alum  is  used  very  extensively  in  the 
preparation  of  pigments,  as  a  mordant,  in  the  sizing  of 
paper,  for  clarifying  water,  etc. 

Ammonium  Alum,  Ammonium  -  Aluminium  Sulphate, 
<NH4)A1(SO4)2  +  12H2O,  is  in  every  way  much  like  the 
potassium  compound,  and  can  be  used  in  place  of  it  for 
almost  all  purposes  for  which  alum  is  used.  It  is  some- 
what more  easily  soluble  than  ordinary  alum.  As  it  is 
cheaper  than  the  latter  it  is  largely  manufactured  in 
place  of  it. 

Sodium  Alum  is  much  more  easily  soluble  in  water 
than  either  potassium  or  ammonium  alum,  and  this 
makes  it  difficult  to  prepare  it  in  pure  condition.  It  is 
therefore  not  manufactured,  although  sodium  compounds 
are  cheaper  than  those  of  either  potassium  or  ammonium. 

Aluminium  Silicate. — It  has  been  stated  that  alumin- 
ium silicate  enters  into  the  composition  of  a  number  of 
important  minerals.  It  occurs  in  enormous  quantities 
in  nature.  The  most  important  of  the  minerals  contain- 
ing it  are  the  feldspars,  of  which  ordinary  feldspar, 
KAlSi3O8,  and  albite,  or  sodium  feldspar,  NaAlSi3O8,  are 
the  most  abundant.  These  again  enter  into  the  compo- 
sition of  granite  together  with  quartz  and  mica,  and  mica 
is  itself  a  double  silicate  of  aluminium.  As  remarked 
under  Silicic  Acid  (which  see),  the  natural  silicates  are  for 
the  most  part  salts  of  polysilicic  acids  which  are  derived 
from  orthosilicic  acid  by  loss  of  water  from  two'  or  more 
molecules.  Up  to  the  present  but  little  more  has  been 
done  with  the  many  natural  silicates  than  to  determine 
their  percentage  composition.  It  appears  probable  from 


580  INORGANIC  CHEMISTRY. 

what  lias  already  been  learned  regarding  their  constitu- 
tion that  investigations  in  this  direction  will  before  long 
yield  interesting  results.  As  yet,  however,  the  methods 
for  such  investigations  are  quite  unsatisfactory,  owing 
largely  to  the  fact  that  the  compounds  are  so  extremely 
stable  that  but  few  reagents  decompose  them,  and  if 
they  are  decomposed  at  all,  the  products  are  such  that 
no  conclusion  can  be  drawn  regarding  the  constitution. 
A  careful  study  of  the  relations  in  which  minerals  occur 
in  nature  will  undoubtedly  be  of  assistance,  as  this  will 
throw  some  light  upon  the  conditions  under  which  they 
were  formed. 

One  of  the  most  common  decompositions  of  minerals, 
constantly  taking  place,  and  that  has  taken  place  to  an 
enormous  extent,  is  that  of  feldspar.  Under  the  in- 
fluence of  moisture  and  the  carbon  dioxide  of  the  air, 
this  substance  slowly  decomposes,  the  products  being 
mainly  potassium  or  sodium  silicate  and  aluminium  sili- 
cate. The  salt  of  the  alkali  metal,  principally  potas- 
sium, being  soluble,  is  carried  away,  and  finds  its  way 
into  the  soil.  The  silicate  of  aluminium  is  not  soluble, 
but  it  easily  forms  emulsions  with  water,  and  is  there- 
fore carried  down  the  sides  of  the  hills  and  mountains 
upon  which  it  is  formed  into  the  valleys,  and  much  of  it 
finds  its  way  into  streams.  Sometimes  this  carrying 
away  is  prevented,  and  then  large  beds  of  comparatively 
pure  clay,  known  as  kaoline,  are  formed.  The  clay 
found  in  the  valleys  is  always  more  or  less  impure  and 
colored. 

Kaoline. — This  is  the  purest  form  of  aluminium  sili- 
cate found  in  nature.  It  always  contains  water.  Its 
composition  varies,  some  specimens  on  analysis  giving 
results  which  lead  to  the  formula  Al4(SiO4)3  +  4H2O, 
according  to  which  the  substance  is  the  salt  of  normal 
silicic  acid,  Si(OH)4.  Other  specimens  have  the  compo- 

roH 

sition  HAlSiO4  +  H3O,   or  Si  J  Q\ M  ,  H  Q  .      When 

[o/ 


CLA  T—  ULTRAMARINE.  ,        581 

heated  alone  kaoline  does  not  melt ;  but  if  feldspar  is 
added  to  it,  the  whole  melts,  and  forms  a  translucent 
mass  known  as  porcelain.  Other  substances  besides 
feldspar  may  be  used  for  this  purpose. 

Clay. — Ordinary  clay,  as  has  been  stated,  is  a  name 
given  to  the  impure  varieties  of  aluminium  silicate  which 
have  been  carried  down  from  the  place  of  formation. 
Among  the  substances  besides  aluminium  silicate  found 
in  clays  are  calcium  carbonate,  magnesium  carbonate, 
sand,  and  hydroxides  of  iron.  The  color  is  largely  deter- 
mined by  the  amount  of  the  hydroxides  of  iron  present. 
The  better  varieties  are  used  in  the  manufacture  of  the 
so-called  "  stone-ware,"  gas-retorts,  and  fire-bricks.  The 
colored  varieties  are  used  for  making  ordinary  earthen- 
ware and  bricks.  Marl  is  clay  mixed  with  considerable 
quantities  of  calcium  carbonate. 

Ultramarine. — The  substance  occurring  in  nature  and 
known  as  lapis  lazuli  consists  of  a  silicate  of  sodium  and 
aluminium  together  with  a  sulphur  compound,  probably 
a  polysulphide  of  sodium.  The  coloring  matter,  known 
as  ultramarine,  obtained  by  powdering  it  was  formerly 
very  expensive,  but  it  is  now  made  artificially  by  the  ton, 
and  the  color  of  the  artificially  prepared  substance  is 
even  more  beautiful  than  that  of  the  natural.  A  great 
deal  of  work  has  been  done  in  the  way  of  investigating 
the  chemical  constitution  of  ultramarine,  but  the  problem 
has  not  yet  been  fully  solved.  The  artificial  preparation 
is  effected  by  melting  together  kaoline,  anhydrous  so- 
dium carbonate,  and  sulphur ;  or  clay,  calcined  sodium 
sulphate,  and  charcoal.  By  varying  the  conditions  of 
the  preparation  products  of  different  colors  are  obtained. 
Besides  the  deep-blue  ultramarine,  there  are  now  manu- 
factured ultramarines  of  different  shades  of  blue,  and  a 
green  variety.  A  white,  a  red,  a  yellow,  and  a  violet, 
variety  are  also  known.  The  substance  which  gives  to 
ultramarine  its  color  is  destroyed  by  acids,  but  not  by 
alkalies.  It  can  be  heated  in  a  closed  vessel  without 
change,  but  if  heated  to  a  high  temperature  in  the  air  or 
in  oxygen  the  color  is  destroyed. 


582  INORGANIC  CHEMISTRY. 

Ultramarine  is  now  manufactured  in  very  large  quan- 
tity— according  to  a  recent  report,  to  the  extent  of  nearly 
9000  tons  a  year.  It  is  the  most  extensively  used  blue 
coloring  matter. 

Porcelain. — It  was  stated  above  that  when  kaoline  is 
heated  alone  it  does  not  melt,  but  that  if  feldspar  is 
added  to  it,  or  if  that  found  in  nature  contains  feldspar, 
as  is  frequently  the  case,  it  either  fuses  together  forming 
a  compact  mass,  or  it  melts  and  forms  a  translucent 
mass.  Further,  when  kaoline  or  any  other  variety  of 
clay  is  mixed  with  water,  a  plastic  substance  results, 
which  can  be  kneaded  and  worked  into  any  desired 
form.  These  facts  form  the  basis  of  the  manufacture  of 
earthenware,  porcelain,  etc.  The  ease  with  which  the 
mass  melts  depends  upon  the  quantity  of  feldspar  or 
other  flux  added  to  it.  If  but  little  is  added  it  melts 
with  difficulty  ;  if  much  is  added  it  melts  easily. 

In  the  manufacture  of  the  finest  kinds  of  porcelain 
kaoline  is  used.  This  is  generally  mixed  with  a  little 
feldspar  or  chalk,  gypsum  or  some  other  flux,  and  sand 
is  also  added.  All  these  substances  must  be  very  finely 
ground.  The  mixture  is  then  worked  into  the  desired 
forms,  and  carefully  dried.  After  the  objects  are  dried 
they  are  next  burned,  first  at  a  red  heat  at  which  the 
mass  becomes  solid,  afterwards  at  a  white  heat  for  the 
purpose  of  forming  a  glaze  upon  the  surface.  The  prod- 
uct after  the  first  burning  is  that  which  is  familiar  as 
porous  earthenware ;  that  formed  in  the  second  burning 
is  the  porcelain  with  glaze  as  it  is  commonly  used. 

In  order  to  form  the  glaze  upon  the  porcelain  the 
porous  earthenware  first  formed  is  drawn  through  a 
vessel  containing  proper  materials  in  finely  powdered 
condition  and  suspended  in  water.  The  materials  used 
are  generally  the  same  as  those  used  for  the  porcelain 
itself,  but  they  are  mixed  in  different  proportions,  witL 
less  kaolin,  and  more  sand  and  feldspar,  so  as  to  be 
more  easily  fusible.  After  this  treatment  the  objects 
are  again  heated  to  a  high  temperature. 

Earthenware. — The  ordinary  varieties  of  earthenware 
are  made  from  varieties  of  clay  which  are  much  less  pure 


REACTIONS  OF  ALUMINIUM  SALTS.  583 

than  kaoline.  Ordinary  colored  clay  is  used.  The  ob- 
jects are  formed,  and  then  subjected  in  general  to  the 
same  kind  of  treatment  as  porcelain.  They  are  glazed 
in  different  ways.  One  method  consists  in  bringing  the 
glazing  material  on  the  earthenware  before  it  is  burned ; 
another  method  consists  in  putting  the  objects  in  the 
furnace  without  a  glaze,  and  towards  the  end  of  the 
firing  process  sodium  chloride  is  thrown  into  the  fur- 
nace, and  is  thus  brought  in  contact  with  the  ware  in 
the  form  of  vapor.  A  chemical  change  takes  place,  re- 
sulting in  the  formation  of  a  silicate  of  aluminium  and 
sodium  upon  the  surface.  This  melts,  and  forms  a  glaze. 

Bricks  are  the  most  common  variety  of  unglazed 
earthenware.  Owing  to  the  presence  of  other  sub- 
stances besides  aluminium  silicate,  as,  for  example,  cal- 
cium carbonate,  the  material  is  comparatively  easily 
fusible.  The  color  of  bricks  is  largely  due  to  the  pres- 
ence of  oxides  of  iron. 

Reactions  of  Aluminium  Salts  which  are  of  Special 
Value  in  Chemical  Analysis. — Potassium  and  sodium  hy- 
droxides precipitate  aluminium  hydroxide,  which  is  solu- 
ble in  an  excess  of  the  reagents. 

Ammonia  precipitates  the  hydroxide,  which  is  only 
slightly  soluble  in  an  excess  of  the  reagent. 

Hydrogen  sulphide  and  carbon  dioxide  precipitate  alu- 
minium hydroxide  from  a  solution  of  an  aluminate ;  that 
is,  from  a  solution  of  aluminium  hydroxide  in  a  caustic 
alkali. 

Ammonium  sulphide  and  other  soluble  sulphides  pre- 
cipitate the  hydroxide.  This  is  due  to  the  instability  of 
the  sulphide  of  aluminium,  or,  going  farther  back,  to 
the  weak  basic  character  of  the  hydroxide.  The  reaction 
of  ammonium  sulphide  with  aluminium  sulphate  takes 
place  as  represented  in  the  following  equation : 

A12(SO4)3  +  3(NH4)2S  +  6H2O  =  3(NH4)2S04  +  3H2S  +  2A1(OH)8. 

Soluble  carbonates  precipitate  aluminium  hydroxide  for 
the  same  reason  that  the  soluble  sulphides  do.  The  re- 


584  INORGANIC  CHEMISTRY. 

action  between  aluminium  sulphate  and  sodium  carbon- 
ate takes  place  thus  : 

Al3(S04)3+3Na2C08+3H20=3NaaSO4+3CO2+2Al(OH),. 

OTHEK  MEMBERS  OF  FAMILY  III,  GROUP  A. 

Scandium,  Sc  (At.  Wt.  43.78). — This  element  was  dis- 
covered in  1880  in  the  minerals  euxenite  and  gadolinite. 
Its  compounds  are  similar  to  those  of  aluminium.  It 
forms  an  oxide  of  the  formula  Sc2O3 ;  a  sulphate,  Se2(SO4)3 ; 
a  double  sulphate,  KSc(SO4)  2;  etc.  It  is  of  special  in- 
terest for  the  reason  that  its  properties  were  foretold  by 
Mendeleeff  several  years  before  it  was  discovered.  The 
prophecy  was  based  upon  the  position  of  the  element  in 
the  periodic  system.  When  the  relations  between  the 
atomic  weights  and  properties  of  the  elements  were  first 
described  in  a  comprehensive  way  by  Lothar  Meyer  and 
Mendeleeff,  the  latter  described  the  properties  of  an  ele- 
ment then  unknown,  which  he  called  ekaboron.  This 
should  have  the  atomic  weight  about  44,  should  form  an 
oxide  of  the  formula  M2O3 ,  etc.  It  has  been  shown  that 
the  properties  of  scandium  agree  very  closely  with  those 
foretold. 

Yttrium,  Y  (At.  Wt.  88.35),  like  scandium,  is  found  in 
gadolinite,  euxenite,  and  some  other  rare  minerals. 
The  element  itself  has  not  been  isolated.  Its  chloride, 
YC13  -\-  6H2O,  is  easily  made.  With  potassium  and  so- 
dium chlorides  it  forms  double  chlorides  analogous  to 
those  formed  by  aluminium  chloride.  Its  oxide  has  the 
formula,  T2O3,  and  is  formed  by  heating  the  hydroxide, 
Y(OH)3,  or  nitrate,  Y(NO3)3.  The  hydroxide,  Y(OH)3,  is 
precipitated  by  adding  potassium  hydroxide  to  a  solu- 
tion of  an  yttrium  salt,  and  is  not  dissolved  by  an  excess 
of  the  alkali.  The  hydroxide,  while  being  less  acidic 
than  aluminium  hydroxide,  is  also  more  strongly  basic, 
as  is  shown  by  its  power  to  unite  with  weak  acids. 
When  exposed  to  the  air,  it  attracts  and  combines  with 
carbonic  acid. 


THE  BORON-ALUMINIUM  GROUP  IN  GENERAL.    585 

Ytterbium,  Yb  (At.  Wt.  171.88). — This  rare  element,  like 
scandium  and  yttrium,  is  found  in  gadolinite  and  euxen- 
ite — most  abundantly  in  the  latter.  Its  compounds  in 
general  resemble  those  of  yttrium.  Its  hydroxide  is  not 
soluble  in  alkalies,  but  it  absorbs  and  combines  with  car- 
bon dioxide.  Its  oxide  has  the  formula  Yb2O3,  its  sul- 
phate, Yb2(SO4)3,  etc. 

Samarium,  Sm,  Terbium,  Tr,  and  Gadolinium,  Gd,  are 
elements  that  have  been  obtained  from  samarskite,  a 
North  Carolina  mineral.  They  have  also  been  found  in 
small  quantities  in  some  other  minerals,  as  cerite,  gado- 
linite,  and  orthite.  In  general  these  elements  resemble 
aluminium  in  their  chemical  conduct. 

Helium. — When  cleveite  and  some  other  rare  minerals 
are  treated  with  acids  a  gas  is  given  off.  This  was  at 
first  held  to  be  nitrogen,  but  a  careful  spectroscopic 
examination  by  W.  Ramsay  has  shown  that  it  contains 
a  hitherto  unknown  gas.  The  spectrum  contains  lines 
which  are  identical  with  lines  that  have  been  observed 
in  the  solar  spectrum.  These  have  been  ascribed  to  an 
unknown  element,  helium.  This  has  now  been  isolated 
and  studied.  It  proves  to  be  a  very  light  gas,  as  incapa- 
ble of  forming  chemical  compounds  as  argon.  Practically 
all  the  evidence  goes  to  show  that  it  is  an  element,  but 
as  yet  little  is  known  in  regard  to  its  relations  to  other 
elements. 

The  chemistry  of  lanthanum  is  so  intimately  connected 
with  that  of  cerium  and  didymium,  that,  although  these 
three  elements  appear  to  belong  to  different  families, 
they  will  be  briefly  considered  together. 

The  Boron- Aluminium  Group  in  General. — Comparing 
the  group  of  which  boron  and  aluminium  are  the  prin- 
cipal members  with  the  potassium,  calcium,  and  mag- 
nesium groups,  it  will  be  seen  that  the  members  of  this 
group  do  not  resemble  one  another  as  closely  as  the 
members  of  the  other  groups  do.  There  is,  however,  the 
same  strengthening  of  the  basic  properties  and  weaken- 
ing of  the  acid  properties  as  the  atomic  weight  increases. 
Boron  is  the  most  strongly  acid  and  the  least  basic ; 
-aluminium  is  more  basic,  but  has  still  some  acid  prop- 


586 


INORGANIC  CHEMISTRY. 


erties ;  while  the  other  members  are  more  strongly 
basic,  and  do  not  exhibit  any  acid  properties.  Compar- 
ing the  first  members  of  Group  A,  of  Families  I,  II,  and 
III,  it  is  clear  that  with  increasing  atomic  weight  the 
acid  properties  and  the  valence  increase.  The  elements 
referred  to  are  lithium,  glucinum,  and  boron. 

Members  of  Group  B,  Families  I,  II,  III,  and  IV.— The 
base-forming  elements  thus  far  considered  form  the  prin- 
cipal groups  of  the  first  three  families.  In  these  principal 
groups  the  most  characteristic  elements  of  these  families 
occur.  But  besides  the  principal  group  of  each  family 
there  is  a  secondary  group,  the  members  of  which  differ 
in  some  respects  from  those  of  the  principal  group, 
though  they  resemble  one  another.  Between  the  second- 
ary group  of  Family  I  and  that  of  Family  II,  further, 
there  are  some  points  of  resemblance.  The  secondary 
group  in  Family  IV  bears  to  the  principal  group  much 
the  same  relation  that  the  secondary  groups  of  the  first 
three  families  bear  to  the  principal  groups.  In  the  table 
p.  151  the  principal  groups  are  those  which  fall  under 
the  letter  A,  and  the  secondary  groups  are  those  which 
fall  under  the  letter  B  in  each  of  the  first  four  families. 
These  secondary  groups  are  : 


Family 

Group 

I 

B 

Copper 

Silver 

Gold 

II 

B 

Zinc 

Cadmium 

Mercury 

III 

B 

Gallium 

Indium 

Thallium 

IV 

B 

Germanium 

Tin 

Lead 

In  the  fifth,  sixth,  and  seventh  families  the  most  char- 
acteristic elements  are  those  which  occur  in  Group  B. 
These  have  already  been  studied. 

Haying  thus  considered  the  members  of  the  principal 
groups  of  the  first  four  families,  let  us  next  turn  to  the 
study  of  the  members  of  the  secondary  groups. 


CHAPTER  XXVIII. 

ELEMENTS  OF  FAMILY  I,  GROUP  B : 
COPPER— SILVER— GOLD. 

General. — The  facts  which  strike  one  most  forcibly  on 
comparing  the  elements  of  this  group  with  those  of 
Group  A  of  the  same  family  are,  that  they  are  much  less 
active  chemically,  and  that  they  furnish  a  greater  variety 
of  compounds.  Sodium  and  potassium  and  the  other 
members  of  Group  A  display  the  greatest  activity,  as  we 
have  seen.  The  basic  character  is  most  strongly  devel- 
oped in  them.  Further,  in  nearly  all  their  compounds 
they  act  with  the  same  valence.  They  are  univalent  in 
all  their  salts.  Copper,  silver,  and  gold,  however,  are 
not  chemically  active  elements,  and  the  activity  grows 
less  with  increasing  atomic  weight.  Copper  and  gold 
form  two  series  of  compounds  each,  and  silver  also  forms 
a  few  compounds,  in  which  it  appears  with  a  valence 
greater  than  one.  In  the  two  series  of  salts  formed  by 
copper  the  element  appears  to  be  univalent  and  bivalent, 
as  in  the  chlorides  CuCl  and  CuCl2.  Gold,  however,  is 
univalent  and  trivaleiit,  while  silver  is  almost  exclusively 
univalent.  It  must  be  said  that  the  resemblance  between 
gold  and  the  other  members  of  Group  B  is  apparently 
not  as  marked  as  that  between  mercury  and  copper  and 
silver.  It  is,  however,  possible  that  as  investigation 
proceeds  the  resemblance  will  appear  more  striking  than 
it  does  at  present. 

COPPER,  Cu  (At.  Wt.  63.12). 

General. — The  compounds  of  .copper  which  are  most 
commonly  met  with  are  those  in  which  it  acts  as  a  biva- 
lent element.  Its  principal  compounds  are  copper  oxide, 

(587) 


588  INORGANIC  CHEMISTRY. 

CuO  ;  the  sulphate,  CuSO4  ;  and  the  sulphide,  CuS.  In 
all  these  the  copper  is  bivalent.  But  besides  these  there 
are  such  compounds  as  CuCl  and  Cu2O,  in  which  the 
element  appears  to  be  univalent.  There  are,  then,  two 
series  of  salts,  of  which  the  following  will  serve  as 
examples  : 

CuCl  CuCla 

CuBr  CuBr, 

Cu20  CuO 

Those  compounds  which  are  of  the  first  order,  corre- 
sponding to  the  chloride  CuCl,  are  called  cuprous  com- 
pounds. Thus,  CuCl  is  cuprous  chloride;  Cu2O,  cuprous 
oxide,  etc.  On  the  other  hand,  compounds  of  the  second 
order  are  called  cupric  compounds.  Thus,  CuCl2  is  cupric 
chloride;  CuO,  cupric  oxide;  CuSO4,  cupric  sulphate,  etc. 
It  has  been  suggested  that  perhaps  the  formula  of  the 
simpler  cuprous  compounds,  like  CuCl,  etc.,  should  be 
doubled,  and  written  Cu2Cl2,  Cu2I2,  etc.  This  suggestion 
has  its  origin  in  the  valence  hypothesis.  In  cupric 
chloride,  CuCl2,  and  cupric  oxide,  CuO,  copper  is  evi- 
dently bivalent  ;  whereas,  if  the  formulas  of  the  cuprous 
compounds  are  the  simpler  ones,  CuCl,  Cul,  etc.,  copper 
is  univalent  in  them.  If,  however,  cuprous  chloride  is 
Cu2Cl2,  it  may  be  that  in  it  the  copper  is  bivalent.  It  is 
only  necessary  to  assume  that  in  the  molecule  of  cu- 
prous chloride  two  atoms  of  copper  are  combined  as 
represented  thus  : 

Cu- 


If,  then,  each  of  the  copper  atoms  should  combine  with 
a  chlorine  atom,  the  compound  would  have  the  formula 
Cu2Cl2.  The  question  here  presented  is  similar  to  that 
concerning  the  molecular  formula  of  aluminium  chloride. 
A  determination  of  the  specific  gravity  of  the  vapor  of 
cuprous  chloride  has  been  made,  and  it  has  been  found 
to  correspond  to  that  required  by  the  formula  Cu2Cl2. 
It  is  possible,  however,  that  at  a  higher  temperature  a 
different  result  may  be  obtained,  as  in  the  case  of  alu- 


COPPER.  589 

minium  chloride,  and  it  is  possible  that  the  compound 
may  be  a  double  chloride,  formed  by  union  through  the 
chlorine  atoms  as  represented  in  the  formula  CuCl-ClOu 
or  Cu-(Cl2)-Cu.  Then  there  would  be  complete  analogy 
between  cuprous  chloride  and  cuprous  oxide,  Cu-O-Cu. 
Our  knowledge  in  regard  to  this  matter  is  extremely 
limited  at  present,  and  it  seems  perfectly  justifiable  to 
use  the  simpler  formulas  for  the  cuprous  compounds 
until  further  evidence  has  been  produced.  Whatever 
the  explanation  may  be,  it  is  undoubtedly  a  fact  that 
there  are  two  series  of  salts  of  copper,  in  one  of  which 
there  is  relatively  half  as  much  copper  as  in  the  other, 
and  it  is  also  a  fact  that  by  comparatively  simple  methods 
the  salts  of  one  series  can  be  converted  into  those  of  the 
other,  as  will  be  pointed  out  below. 

Forms  in  which  Copper  occurs  in  Nature. — Copper  is 
a  widely  distributed  element,  and  it  occurs  also  in  large 
quantities.  It  occurs  in  the  uncombined  condition,  or 
as  native  copper,  in  large  quantity  in  the  United  States 
in  the  neighborhood  of  Lake  Superior,  in  China,  Japan, 
Siberia,  and  Sweden.  The  most  valuable  ores  of  cop- 
per are  the  oxides,  ruby  copper  or  cuprous  oxide,  Cu2O, 
and  cupric  oxide,  CuO ;  the  carbonates,  as  malachite, 
Cu2(OH)2CO3 ;  the  sulphides,  as  chalcocite,  Cu2S,  copper 
pyrites,  Cu2S.Fe2S3 ;  and  others. 

Metallurgy  of  Copper. — The  metallurgy  of  copper  is 
comparatively  complicated,  owing  to  the  difficulty  of 
converting  the  ores  of  copper  into  the  oxide.  In  most 
of  the  ores  used  sulphur  and  iron  are  contained,  as  well 
as  smaller  quantities  of  other  elements,  as  arsenic,  anti- 
mony, lead,  etc.  The  ores  are  first  roasted  with  the 
object  of  converting  the  sulphides  partly  into  oxides. 
Under  these  circumstances  the  sulphides  of  iron  are 
more  easily  converted  into  the  oxides  than  the  sulphides 
of  copper.  By  adding  a  material  rich  in  silicic  acid,  and 
melting  the  roasted  ore  in  a  blast  furnace  with  charcoal, 
the  oxide  of  iron  is  partly  reduced,  and  converted  into 
silicate,  which  runs  off  with  the  slag.  In  this  way  a 
product  is  obtained  which  is  richer  in  copper  than  the 
roasted  ore.  This,  which  is  called  the  matte,  contains 


590  INORGANIC  CHEMISTRY. 

copper  sulphide  and  iron  sulphide.  The  matte  is  again 
roasted  and  melted  in  the  same  way  as  the  ore,  and  a 
further  quantity  of  iron  is  removed,  while*  some  of  the 
copper  is  reduced.  A  reaction  which  plays  an  impor- 
tant part  in  these  processes  is  that  which  takes  place 
between  cuprous  oxide  and  cuprous  sulphide,  forming 
metallic  copper  and  sulphur  dioxide : 

2Cu2O  +  Cu2S  =  6Cu  +  S02. 

Sometimes  it  is  necessary  to  repeat  the  roasting  and 
melting  with  charcoal  and  sand  a  number  of  times,  the 
matte  becoming  richer  in  copper  at  each  successive  stage. 
Properties. — Copper  is  a  hard  metal,  of  a  reddish  color 
and  metallic  lustre.  It  does  not  change  in  dry  air,  but 
in  moist  air  it  gradually  becomes  covered  with  a  green 
layer  of  a  basic  carbonate.  It  melts  at  a  somewhat 
lower  temperature  than  gold,  and  at  a  somewhat  higher 
temperature  than  silver.  It  is  very  malleable  and  tena- 
cious. It  decomposes  water  only  at  bright-red  heat. 
When  heated  in  the  air  to  a  comparatively  high  tempera- 
ture it  becomes  covered  with  a  layer  of  cupric  oxide  ;  at 
a  lower  temperature  cuprous  oxide  is  formed.  Nitric 
acid  dissolves  it,  copper  nitrate,  Cu(NO3)2,  being  formed, 
and  the  oxides  of  nitrogen  being  evolved  (see  p.  285) ; 
dilute  sulphuric  acid  does  not  act  upon  it  unless  the  air 
has  access  to  it ;  concentrated  sulphuric  acid  when 
heated  with  it  forms  cupric  sulphate,  CuSO4,  and  sul- 
phur dioxide  (see  p.  217).  Dilute  acids  in  general  do 
not  act  upon  it  unless  the  air  has  access  to  it.  This  fact 
is  of  importance  in  connection  with  the  use  of  copper 
vessels  in  culinary  operations-  Substances  containing 
vegetable  acids  can  be  boiled  in  bright  copper  vessels 
with  impunity,  for  the  water  vapor  prevents  the  access  of 
air,  but,  on  cooling,  the  air  is  admitted,  and  then  action 
takes  place,  causing  solution  of  some  of  the  copper,  which 
is  objectionable.  Ammonia  in  contact  with  copper  ab- 
sorbs oxygen,  and  the  copper  dissolves  in  consequence 
of  the  formation  of  a  compound  of  cupric  oxide  and  am- 
monia. This  fact  is  sometimes  taken  advantage  of  for 


ALLOTS  OF  COPPER,  591 

the  preparation  of  nitrogen,  as  was  stated  in  speaking  of 
this  gas  (see  p.  249). 

Applications. — As  is  well  known,  copper  is  used  very 
extensively  for  a  variety  of  purposes,  among  which  the 
following  may  be  mentioned :  for  electrical  apparatus, 
coins,  copper  vessels,  roofs,  for  covering  the  bottoms  of 
ships,  etc.  It  is  also  used  for  copper-plating  ;  and  in  the 
preparation  of  a  number  of  valuable  alloys,  such  as 
brass,  bronze,  gun-metal,  bell-metal,  etc. 

Alloys. — Brass  is  a  mixture  or  compound  of  about  one 
part  of  zinc  and  two  parts  of  copper ;  these  proportions 
may,  however,  be  varied  between  quite  wide  limits. 
There  is  a  variety  of  brass  containing  equal  parts  of 
zinc  and  copper,  and  another  containing  one  part  of  zinc 
and  five  parts  of  copper.  Pinchbeck  is  made  by  combin- 
ing two  parts  of  copper  and  one  of  brass. 

Bronze  consists  of  copper,  zinc,  and  tin.  The  propor- 
tion of  copper  varies  from  65  to  84  per  cent ;  that  of 
zinc  from  31.5  to  11  per  cent ;  and  that  of  tin  from  2.5 
to  4  per  cent.  When  exposed  to  the  air  bronze  becomes 
covered  with  a  green  coating  of  basic  copper  carbonate, 
which  protects  it  from  further  action.  This  coating  is 
now  generally  produced  artificially  by  a  variety  of  meth- 
ods, as  by  washing  the  surface  with  a  solution  of  salts 
and  acids. 

Gun-metal  consists  generally  of  copper  and  tin  in  the 
proportion  of  11  parts  of  tin  and  100  parts  of  copper. 

Bell-metal  contains  a  larger  proportion  (from  20  to  25 
per  cent)  of  tin  than  gun-metal. 

Alloys  with  Aluminium  containing  aluminium  and  cop- 
per in  widely  different  proportions  are  made.  That  with 
3  per  cent  of  copper  has  a  whiter  color  than  aluminium, 
the  color  being  more  like  that  of  silver.  On  the  other 
hand,  an  alloy  of  copper  with  5  to  10  per  cent  of  alu-. 
minium  has  a  color  resembling  that  of  gold.  This,  which 
is  known  as  aluminium  bronze,  is  very  hard  and  elastic, 
and  is  not  easily  acted  upon  by  chemical  reagents.  It  is 
now  used  to  a  considerable  extent  in  the  manufacture  of 
ornamental  and  useful  articles. 

German  silver  is  an  alloy  consisting  of  copper,  zinc,  and 


592  INORGANIC  CHEMISTRY. 

nickel.  The  proportion  of  copper  varies  from  40  to  60 
per  cent ;  that  of  zinc  from  19  to  44  per  cent ;  and  that 
of  nickel  from  6  to  22  per  cent. 

Cuprous  Hydride,  CuH. — This  compound  is  made  by 
treating  a  solution  of  barium  hypophosphite  with  ,a  solu- 
tion of  copper  sulphate.  It  is  thrown  down  as  a  yellow 
precipitate  which  gradually  becomes  darker.  At  60°  it 
decomposes  into  copper  and  hydrogen.  With  hydro- 
chloric acid  it  yields  cuprous  chloride  and  hydrogen. 

Cupric  Hydride,  CuH2,  is  formed  by  treating  a  solution 
of  copper  sulphate  with  hypophosphorous  acid.  When 
freshly  prepared  it  is  a  reddish-brown  sponge-like  mass, 
which,  however,  changes  to  a  chocolate-colored  powder 
on  being  freed  from  acid  and  boiled  for  some  time.  It 
is  not  readily  changed  when  heated  in  the  air.  It  dis- 
solves  in  hydrochloric  acid  with  evolution  of  hydrogen. 

Cuprous  Chloride,  CuCl,  is  formed  by  heating  cupric 
chloride,  CuCl2 ;  by  passing  hydrochloric  acid  over  highly 
heated  copper ;  by  boiling  an  excess  of  copper  filings  with 
concentrated  hydrochloric  acid  with  the  addition  of  a 
little  nitric  acid,  filtering  through  asbestos,  and  pouring 
into  water  ;  and  by  treating  cupric  chloride  with  reducing 
agents,  as,  for  example,  sulphurous  acid.  The  action  with 
sulphur  dioxide  takes  place  as  represented  in  the  equation 

2CuCla  +  SO3  +  2H20  =  2CuCl  +  H2SO4  +.2HC1. 

It  is  a  white  crystalline  compound,  and  is  difficultly  solu- 
ble in  water.  When  exposed  to  the  air  it  rapidly  turns 
green  in  consequence  of  the  formation  of  a  basic  chloride, 
as,  for  example,  2CuO.CuCl2,  or  Cl-Cu-O-Cu-O-Cu-Cl. 
It  is  volatile  at  a  high  temperature,  and  a  determination 
of  the  specific  gravity  of  the  vapor  gave  a  result  corre- 
sponding to  the  formula  Cu2Cl2,  as  was  stated  above.  It 
has  markedly  the  power  to  absorb  chlorine,  and  there- 
fore acts  as  a  reducing  agent.  Ammonia  dissolves  itr 
forming  a  compound  of  the  composition  represented  by 
the  formula  CuCl.NH3,  which  may  be  regarded  as  de- 
rived from  ammonium  chloride  by  replacing  an  atom  of 
hydrogen  by  an  atom  of  copper. 

Cupric  Chloride,  CuCl2. — This  compound  is  formed  by 
treating  copper  or  cuprous  chloride  with  chlorine.  It  is 


COMPOUNDS  OF  COPPER.  593 

also  easily  made  by  dissolving  cupric  hydroxide,  or  car- 
bonate, in  hydrochloric  acid.  From  its  solution  in  water 
the  chloride  crystallizes  with  two  molecules  of  water, 
CuCla  -|-  2H2O.  The  crystals  when  heated  lose  their 
water  without  suffering  further  decomposition,  except  at 
high  heat,  when  a  part  of  the  chlorine  is  given  off,  and 
cuprous  chloride  is  formed.  Cupric  chloride  combines 
with  ammonia  gas,  forming  a  compound  of  the  formula 
CuCl2.6NH3,  which  is  soluble  in  water,  with  a  dark  blue 
color.  When  heated  it  loses  four  molecules  of  ammonia, 
and  the  compound  CuCl2.2NH3  is  left  behind.  This  may 
be  regarded  as  ammonium  chloride,  in  which  two  hydro- 
gen atoms  have  been  replaced  by  an  atom  of  bivalent 

copper,   as   represented    in    the    formula    ^u<]»jTi3ni- 

There  is,  however,  no  direct  experimental  evidence  in 
favor  of  this  view. 

With  other  chlorides  cupric  chloride  forms  double 
chlorides  similar  to  those  formed  by  magnesium  and 
aluminium ;  such,  for  example,  as  CuCl2.2NH4Cl  or 
(NH4)2CuCl4,  CuCl2.2KCl,  or  K2CuCl4,  etc. 

Cuprous  Iodide,  Cul. — When  a  solution  of  a  cupric  salt 
is  treated  with  potassium  iodide,  cuprous  iodide  is  pre- 
cipitated and  iodine  is  set  free,  owing  to  the  instability 
of  cupric  iodide  : 

2CuS04  +  4KI  =  2K2SO4  +  2CuI  +  I2. 

If  a  reducing  agent  is  added  at  the  same  time,  iodine  is 
not  set  free.  Thus,  for  example,  when  sulphur  dioxide 
is  used,  the  reaction  takes  place  as  represented  in  the 
equation 

2CuSO4+  SO2+  2KI  +  2H2O  =  K2SO4+  2H2SO4+  2CuL 

It  forms  a  white  precipitate.  Cupric  iodide  is  not 
known.  A  similar  conduct  is  shown  by  the  cyanides. 

Cuprous  Hydroxide,  Cu(OH). — The  simple  compound 
of  the  formula  here  given  is  not  known,  but  a  derivative 
of  this,  of  the  formula  Cu8O3(OH)2  or  4Cu2O.H2O,  is 
easily  made  by  adding  sodium  hydroxide  to  a  solution  of 


594  INORGANIC  CHEMISTRY. 

a  cuprous  salt.    It  passes  readily  over  into  cuprous  oxide 
by  gently  heating  it. 

Cuprous  Oxide,  Cu2O,  occurs  in  nature,  and  is  known  as 
r'uhy  copper  or  cuprite.  It  is  easily  prepared  by  treating 
a  solution  of  glucose,  or  starch  sugar,  with  copper  sul- 
phate and  potassium  hydroxide.  By  boiling,  the  copper 
is  thrown  down  in  the  form  of  cuprous  oxide.  At  first 
this  is  yellow,  and  it  is  supposed  by  some  that  the  yel- 
low compound  is  the  hydroxide,  but  satisfactory  evi- 
dence of  the  correctness  of  this  view  has  not  been  fur- 
nished. The  yellow  precipitate  is  soon  converted  into  the 
red  oxide.  Cuprous  oxide  is  not  changed  when  allowed 
to  lie  in  contact  with  the  air.  It  dissolves  in  nitric  and 
sulphuric  acids,  forming  cupric  salts ;  and  if  the  acids 
are  dilute,  copper  is  deposited.  This  will  be  clear  from 
a  consideration  of  the  following  equation  : 

Cu20  +  H2S04  =  CuSO4  +  H2O  +  Cu. 

Cupric  Hydroxide,  Cu(OH)2,  like  the  hydroxides  of  most 
base-forming  elements,  is  thrown  down  by  the  addition  of 
a  soluble  hydroxide  to  a  cupric  salt.  It  is  a  voluminous, 
blue  precipitate.  When  allowed  to  stand  in  a  solution,  or 
when  the  solution  is  boiled,  the  hydroxide  loses  a  part  of 
its  hydroxyl,  and  is  converted  into  a  black  compound  of 
the  formula 

Cu(OH)a  +  2CuO,  or  HO-Cu-O-Cu-O-Cu-OH, 

and  this  when  dried  and  heated  is  converted  into  the 
oxide  CuO. 

Cupric  Oxide,  CuO. — Cupric  oxide  is  found  in  nature  in 
tjie  neighborhood  of  Lake  Superior  in  the  United  States, 
and  is  formed  by  heating  copper  to  redness  in  contact 
with  the  air,  or  by  heating  the  nitrate.  It  loses  its  oxy- 
gen very  readily  when  treated  with  reducing  agents,  such 
as  hydrogen  and  carbon.  It  is  used  extensively  in 
quantitative  analysis  for  the  purpose  of  estimating  the 
composition  of  organic  compounds,  or  such  as  contain 
carbon  and  hydrogen.  Its  use  is  based  upon  the  fact 
that  when  organic  compounds  are  heated  with  the  oxide 


CUPRIC  SULPHATE,  595 

they  are  oxidized,  the  carbon  being  converted  into  car- 
bon dioxide  and  the  hydrogen  into  water.  By  passing 
the  products  of  the  oxidation  through  calcium  chloride, 
and  a  solution  of  potassium  hydroxide,  the  water  is  re- 
tained in  the  first,  and  the  carbon  dioxide  in  the  second, 
and  the  weight  of  each  formed  can  easily  be  determined. 
Cupric  oxide  is  dissolved  by  ammonia  in  the  presence  of 
air  and  a  little  of  some  ammonium  salt.  The  composition 
of  the  compound  in  the  solution  is  not  known. 

Other  Oxides  of  Copper. — Besides  cuprous  and  cupric 
oxides,  copper  forms  two  other  -compounds  with  oxygen. 
These  are  copper  siiboxide,  Cu4O,  and  the  peroxide,  CuO2. 
The  former  is  prepared  by  treating  a  solution  of  copper 
sulphate  with  stannous  chloride.  It  takes  up  oxygen 
from  the  air,  and  is  converted  into  higher  oxides.  The 
peroxide  is  said  to  be  formed  by  treating  cupric  hydrox- 
ide with  hydrogen  peroxide. 

Cupric  Sulphate,  CuSO4. — This  salt  is  manufactured  on 
the  large  scale,  and  in  the  crystallized  form,  containing 
five  molecules  of  water,  CuSO4  -f-  5H2O,  is  commonly 
called  "  blue  vitriol."  It  is  found  in  nature  to  some  ex- 
tent, being  formed  by  the  action  of  the  oxygen  of  the  air 
on  the  sulphide.  It  is  most  conveniently  made  by  dis- 
solving metallic  copper  in  concentrated  sulphuric  acid, 
or  by  treating  cupric  sulphide  with  sulphuric  acid.  The 
action  of  sulphuric  acid  on  the  metal  has  already  been 
referred  to.  It  consists  essentially  in  the  formation  of 
cupric  sulphate,  sulphur  dioxide,  and  water,  as  expressed 
in  the  equation 

Cu  +  2H2S04  =  CuSO4  +  SO2  +  2H2O. 

The  question  whether  the  copper  reduces  the  sulphuric 
acid  directly,  or  the  hydrogen  given  off  from  the  acid 
effects  the  reduction,  is  an  open  one.  But  there  are 
other  products  formed  besides  those  mentioned.  At 
first  a  brown  substance  of  the  composition  Cu2S  is 
deposited.  As  the  action  proceeds  oxysulphides  are 
formed,  the  final  product  of  a  series  of  changes  being 
Cu2OS,  or  CuO.CuS,  which  is  black,  and  insoluble  in 
water.  Under  some  conditions  a  considerable  Droportion 


596  INORGANIC  CHEMISTRY. 

of  the  copper  is  transformed  into  the  oxy  sulphides  by 
sulphuric  acid,  Cupric  sulphate  is  obtained  in  large 
blue  crystals  of  the  triclinic  system,  which  have  the  com- 
position CuSO4  +  5H2O.  When  heated  to  100°,  four 
molecules  of  water  are  given  off,  and  the  last  is  not 
given  off  until  the  temperature  200°  is  reached.  This 
makes  it  appear  probable  that  the  salt  has  the  con- 
stitution represented  by  the  formula  CuSO3(OHj2  or 

[g>Cu 
OS  \  ^TT      ,  corresponding  in  this  respect  to  magnesi- 

[OH 

um  sulphate  (which  see).  When  heated  higher,  it  loses 
all  its  hydroxyl,  and  the  salt,  CuSO4,  is  left  in  the  form 
of  a  white  powder,  which  has  the  power  to  take  up  water 
from  the  air,  becoming  blue  again.  It  dissolves  in  three 
parts  of  cold  water  and  one-half  part  boiling  water.  Cop- 
per sulphate,  containing  seven  molecules  of  water,  CuSO4 
-f-  7H2O,  is  obtained  when  mixed  with  solutions  of  the 
sulphates  of  iron,  zinc,  or  magnesium,  all  of  which  crys- 
tallize with  seven  molecules  of  water.  In  this  form  cu- 
pric  sulphate  is  isomorphous  with  the  other  sulphates. 
These  salts  have  in  general  received  the  name  of  vitriols, 
and  the  old  names  "  green  vitriol,"  "  white  vitriol,"  and 
"blue  vitriol"  are  still  used  to  some  extent,  though 
rarely  by  chemists.  Among  the  similar  salts  included 
under  the  same  general  head  are  the  following : 

Zinc  sulphate  (white  vitriol),    ....  ZnSO4  +  7H2O 

Magnesium  sulphate, MgSO4  -f-  7H2O 

Glucinum  sulphate, G1SO4  +  7H2O 

Ferrous  sulphate  (green  vitriol),  .     .     .  FeSO4  -j-  7H2O 

Nickel  sulphate, MSO4  +  7H2O 

Cobalt  sulphate, CoSO4  +  7H2O 

Copper  sulphate  (blue  vitriol), 

CuSO4  +  7H2O,(CuSO4  +  5H2O) 

Cupric  sulphate  is  used  extensively  in  the  preparation 
of  blue  and  green  pigments,  in  copper-plating  by  elec- 
trolysis, in  galvanic  batteries,  for  the  purpose  of  pre- 
serving wood,  and  as  a  remedy  against  phylloxera  (see 
p.  406),  etc. 


CUPRIC  SULPHATE.  597 

Cupric  sulphate  combines  with  other  sulphates,  form- 
ing double  salts  similar  to  those  formed  by  aluminium 
and  magnesium.  The  potassium  and  ammonium  com- 
pounds have  the  formulas  K2SO4.CuSO4  -f-  6H2O  and 
(NH4)2SO4.CuSO4  +  6H2O,  and  probably  have  the  consti- 
tution represented  by  the  general  formula 

"     so,<§M 

)<>Cu. 


When  a  solution  of  cupric  sulphate  is  treated  with 
ammonia  a  precipitate  of  a  basic  salt  is  at  first  formed, 
but  this  dissolves  when  more  ammonia  is  added,  forming 
a  deep-blue  solution.  The  precipitate  first  formed  is 
a  basic  sulphate  of  copper ;  while  the  solution  contains 
a  compound  of  cupric  sulphate,  ammonia,  and  water,  of 
the  composition  represented  by  the  formula  CuSO4. 
4NH3  -f-  H2O.  When  heated  the  salt  loses  water  and 
ammonia,  until  it  has  the  composition  CuSO4.2NH3. 
This  is  probably  analogous  to  the  ammonia  compound 
of  cupric  chloride,  CuCl2.2NH3,  and  may  be  regarded  as 
having  a  similar  constitution,  that  is,  as  ammonium 
sulphate,  in  which  two  hydrogen  atoms  have  been  re- 
placed by  an  atom  of  bivalent  copper,  as  expressed  in 

the  formula  SO4<^jj3>  Cu.  It  is  a  curious  and  inter- 
esting, though  at  present  inexplicable,  fact,  that  anhy- 
drous copper  sulphate  combines  with  five  molecules  of 
ammonia  just  as  it  does  with  five  molecules  of  water, 
and  that  by  lying  in  moist  air  the  molecules  of  ammonia 
in  the  compound  are  successively  replaced  by  water,  so 
that  the  following  series  of  compounds  is  formed : 

CuSO4.5NH3; 
CuSO4.4NH3.H2O ; 
CuS04.3NH3.2H2O; 
CuSO4.2NH3.3H2O; 
CuS04.NH,.4HaO ; 
CuS04.5H2O, 


598  INORGANIC  CHEMISTRY. 

From  this  it  would  appear  that  the  ammonia  in  these 
compounds  plays  a  part  analogous  to  that  played  by 
the  "  water  of  crystallization."  This  does  not  speak 
in  favor  of  the  view  above  expressed  concerning  the 
constitution  of  cuprammonium  sulphate,  in  which  the 
copper  is  held  to  be  in  combination  with  nitrogen. 
There  is  in  fact  no  satisfactory  theory  for  most  of  the 
salts  containing  water  of  crystallization,  nor  for  most 
of  those  containing  ammonia.  The  power  to  combine 
with  ammonia  is  very  commonly  met  with  among  me- 
tallic salts — probably  fully  as  much  so  as  the  power 
to  combine  with  water.  Some  metals  indeed,  as  cobalt 
and  platinum,  form  a  very  large  number  of  complex 
compounds  with  ammonia,  and  with  ammonium  salts. 

Cupric  Nitrate,  Cu(NO3)2,  is  easily  formed  by  dissolv- 
ing copper  in  dilute  nitric  acid.  It  is  easily  soluble  in 
water,  and  is  deposited  in  crystallized  form,  the  crystals 
containing  three  or  six  molecules  of  water  according  to 
the  temperature,  the  salt  with  six  molecules  being 
formed  at  the  lower  temperature.  Like  other  copper 
salts,  it  has  a  blue  color.  It  combines  with  ammonia 
and  with  ammonium  nitrate. 

Cupric  Arsenite,  CuHAsO3,  is  formed  as  a  greenish- 
yellow  precipitate  when  an  ammoniacal  solution  of  arse- 
nious  acid  is  added  to  a  solution  of  cupric  sulphate.  It 
is  known  as  Scheele's  green.  A  compound  of  cupric  arse- 
nite  and  cupric  acetate,  which  is  made  by  treating  a 
basic  acetate  of  copper  with  arsenious  acid,  is  known  as 
Schweinfurt  green.  On  account  of  their  poisonous  char- 
acter these  compounds  are  not  now  used  as  extensively 
as  formerly. 

Cupric  Carbonates. — When  a  soluble  carbonate  is 
added  to  a  solution  of  cupric  sulphate  a  voluminous 
greenish  precipitate  is  formed,  which  has  the  composition 

Cu<gH 

Cu(OH)2.CuC03,    or  Q>CO.      This   is  plainly  a 

Cu<OH 

basic  carbonate.  The  mineral  malachite,  which  has  a 
beautiful  green  color,  has  the  same  composition  as  the 
precipitate  just  mentioned. 


CYANIDES  OF  COPPER.  599 

Cyanides  of  Copper. — Both  cuprous  and  cupric  cya- 
nides are  known,  but  while  generally  the  cupric  com- 
pounds are  the  more  stable,  cupric  cyanide,  like  cupric 
iodide,  is  extremely  unstable.  It  is  readily  changed  to 
a  compound  intermediate  between  the  cupric  and  the 
cuprous  salt.  This  has  the  composition  CuCy2.2CuCy. 
By  heating  in  suspension  in  water  this  intermediate  com- 
pound is  converted  into  the  cuprous  salt.  The  cuprous 
compound  is  quite  stable.  Cupric  cyanide  is  formed  as  a 
yellow  precipitate,  when  potassium  cyanide  is  added  to 
a  solution  of  a  copper  salt.  It  soon  changes  sponta- 
neously into  the  compound  above  mentioned,  which  has 
a  green  color.  When  this  is  heated  it  yields  cuprous 
cyanide,  CuCN,  which  is  white.  Cuprous  cyanide  is  in- 
soluble in  water.  If  an  excess  of  potassium  cyanide  is 
added  to  a  solution  of  a  copper  salt,  the  precipitate  dis- 
solves in  consequence  of  the  formation  of  double  cya- 
nides similar  to  the  double  chlorides.  One  of  these  has 
the  composition  KCN.CuCN  or  KCu(CN)2.  The  one 
formed  under  ordinary  circumstances  is  SKCN.CuCN  or 
K3Cu(CN)4.  The  double  cyanides  are  in  general  more 
complicated  in  composition  than  the  double  chlorides,  as 
we  shall  see  in  studying  those  which  contain  iron — such, 
for  example,  as  the  salt  already  referred  to  under  the 
name  potassium  ferrocyanide  or  yellow  prussiate  of  pot- 
ash, of  the  composition  K4Fe(CN)6. 

Cuprous  Sulphocyanate,  CuSCN",  and  Cupric  Sulpho- 
cyanate,  Cu(SCN)2,  bear  to  each  other  relations  similar  to 
those  which  exist  between  the  cyanides.  When  potassium 
sulphocyanate  is  added  to  a  concentrated  solution  of  a  cu- 
pric salt  cupric  sulphocyanate  is  precipitated  as  a  black 
powder.  If  the  solution  is  diluted,  decomposition  into  the 
cuprous  salt  takes  place.  When  a  reducing  agent  such  as 
sulphur  dioxide  is  added  at  the  same  time,  only  the 
cuprous  salt  is  formed. '  This  is  a  white,  granular  pow- 
der, insoluble  in  water. 

Cuprous  Sulphide,  Cu2S. — This  compound  occurs  in 
nature,  and  is  known  as  chalcocite.  It  is,  further,  a  con- 
stituent of  copper  pyrites,  which  is  a  compound  of  cu- 
prous and  ferric  sulphides,  CuaS.Fe9S,  or  CuFeS2.  It 


600  INORGANIC  CHEMISTRY. 

can  be  made  by  heating  copper  and  sulphur  together  in 
the  right  proportions.  It  has  a  grayish-black  color; 
does  not  give  up  its  sulphur,  even  when  heated  in  hy- 
drogen ;  and  is  the  more  stable  of  the  two  sulphides  of 
copper. 

Cupric  Sulphide,  CuS. — This  is  formed  as  a  black  pre- 
cipitate when  hydrogen  sulphide  is  passed  into  a 
solution  of  a  cupric  salt.  In  water  alone  cupric  sul- 
phide is  somewhat  soluble.  Hence  in  washing  out  a 
precipitate  of  copper  sulphide  with  water  a  little  of  it 
will  pass  through  in  solution.  It  also  easily  undergoes 
oxid.ation,  and,  as  it  forms  the  sulphate,  some  is  dissolved 
in  this  way  unless  proper  precautions  are  taken.  It  is 
slightly  soluble  in  ammonium  sulphide,  but  insoluble  in 
sodium  sulphide.  The  above  facts  are  of  importance  in 
the  analysis  of  compounds  containing  copper,  as  will 
readily  be  seen.  When  heated,  cupric  sulphide  loses 
half  its  sulphur,  and  is  converted  into  cuprous  sulphide. 

Copper-plating. — The  process  of  copper-plating  con- 
sists in  brief  in  depositing  upon  an  object  a  layer  of  cop- 
per by  putting  it  in  a  bath  containing  some  copper  salt, 
and  connecting  it  with  one  pole  of  an  electric  battery. 
Decomposition  of  the  copper  salt  takes  place,  and  cop- 
per is  deposited  upon  the  object.  Alkaline  solutions  of 
the  double  cyanides  are  best  adapted  to  the  purpose. 
The  process  is  extensively  used  in  the  preparation  of 
electrotype  plates.  These  are  plates  which  are  prepared 
either  from  wood-cuts  or  from  type  by  making  a  mould 
of  gutta-percha,  covering  this  with  graphite,  and  immers- 
ing the  plate  thus  prepared  in  the  copper-plating  bath. 
The  plate  thus  made  is  an  exact  reproduction  of  the 
wood-cut  or  type  of  which  the  impression  in  gutta- 
percha  was  taken. 

Reactions  which  are  of  Special  Value  in  Chemical 
Analysis. — Potassium  or  sodium  hydroxide  forms  a  blue 
precipitate  which  becomes  black  on  standing  or  when 
heated.  (See  Cupric  Hydroxide.) 

Ammonia  first  forms  a  greenish  precipitate,  which  is  a 
basic  salt.  With  cupric  sulphate  the  reaction  takes 
place  thus : 


SILVER.  601 


S02<°>Cu 

+  2NH3  +  2H20  =  >  S02  +  (NH4)aS04. 

S02<o>Cu  Cu<OH 

If  the  action  is  carried  farther,  the  basic  salt  dis- 
solves, forming  the  compound  referred  to  under  Cupric 
Sulphate  (which  see),  the  solution  being  dark  blue. 

Potassium  or  sodium  carbonate  precipitates  the  basic 
carbonate  referred  to  under  Cupric  Carbonate  (which 
see).  The  change  in  color  from  blue  to  green  which 
takes  place  in  this  precipitate  is  probably  due  to  a  loss 
of  water. 

Potassium  ferrocyanide,  K4Fe(CN)6,  forms  a  reddish- 
brown  precipitate,  which  is  the  corresponding  copper 
salt,  Cu2Fe(CN)6.  This  compound  is  decomposed  by 
caustic  alkalies,  forming  cupric  oxide  and  the  corre- 
sponding alkali  salt,  Na4Fe(CN)6  or  K4Fe(CN)6.  The 
reactions  with  potassium  iodide,  cyanide,  and  sulpho- 
cyanide  have  been  explained  above. 

In  the  oxidizing  flame  the  bead  of  borax  or  microcosmic 
salt  is  greenish  blue,  while  when  heated  in  the  reducing 
flame  it  appears  opaque  and  red.  The  red  color  is  due 
to  the  reduction  of  the  oxide  to  copper  or  cuprous  oxide. 

SILVER,  Ag  (At.  Wt.  107.11). 

General.  —  In  nearly  all  the  compounds  of  silver  the 
element  is  univalent.  It,  however,  forms  three  oxides  of 
the  formulas  Ag4O,  Ag2O,  and  AgO.  The  compounds 
correspond  closely  in  many  respects  to  the  cuprous  com- 
pounds. There  is  the  same  question  here  as  in  the  case 
of  copper  as  to  whether  the  molecular  weights  corre- 
spond to  the  simple  formulas  AgCl,  AgBr,  AgNO3,  etc., 
or  to  the  doubled  formulas  Ag2Cl2,  Ag2Br2,  Ag2(NO3)2, 
etc.  There  is  no  evidence  at  present  known  by  which  a 
decision  between  the  two  possibilities  can  be  reached. 
The  simpler  formulas  will  therefore  be  used  here. 

Forms  in  which  Silver  Occurs  in  Nature.  —  Silver  occurs 
to  some  extent  native,  but  for  the  most  part  in  combina- 
tion, particularly  with  sulphur,  and  in  company  with 


602  INORGANIC  CHEMISTRY. 

lead.  The  principal  ore  of  silver  is  the  sulphide,  Ag2S, 
which  occurs  in  combination  with  other  sulphides,  as  of 
lead,  copper,  arsenic,  antimony,  etc.  The  compounds 
with  chlorine,  bromine,  and  iodine  are  also  found,  but 
in  smaller  quantity  than  the  sulphide.  Small  quantities 
of  the  sulphide  are  found  in  almost  all  varieties  of  ga- 
lenite  or  lead  sulphide. 

Metallurgy  of  Silver. — Much  of  the  silver  in  use  is  ob- 
tained from  galenite,  PbS.  This  mineral  is  treated  in 
such  a  way  as  to  cause  the  separation  of  the  lead  (which 
see),  and  the  silver  is  separated  from  sulphur  at  the 
same  time.  But  it  is  dissolved  in  a  large  quantity  of 
lead,  and  the  problem  which  presents  itself  to  the  me- 
tallurgist is  how  to  separate  the  small  quantity  of 
silver  from  the  large  quantity  of  lead.  This  is  accom- 
plished by  melting  the  mixture  and  allowing  it  to  cool 
until  crystals  appear.  These  are  almost  pure  lead. 
They  are  dipped  out  by  means  of  a  sieve-like  ladle,  and 
the  liquid  left  is  again  allowed  to  stand,  when  another 
crop  of  crystals  is  formed,  and  can  be  removed  in  the 
same  way  as  before.  By  this  means,  and  by  again  melt- 
ing the  crystals  removed,  allowing  the  liquid  to  crystal- 
lize, and  removing  the  crystals  formed,  there  is  finally 
obtained  a  product  which  is  rich  in  silver,  but  which  still 
contains  lead.  This  is  heated  in  appropriate  vessels  in 
contact  with  the  air,  when  the  lead  is  oxidized,  while  the 
silver  remains  in  the  metallic  state.  This  method  of 
concentrating  by  crystallization  of  lead  is  known  as  Pat- 
tinsoris  method. 

Another  method  of  separating  lead  and  silver  now  ex- 
tensively used  consists  in  treating  the  molten  alloy  with 
a  small  quantity  of  zinc.  This  takes  up  all  the  silver, 
and  the  alloy  of  zinc  and  silver  thus  formed  is  removed, 
and  afterwards  treated  with  superheated  steam,  by 
which  the  zinc  is  oxidized  and  the  silver  left  unchanged. 

Some  ores  of  silver  are  treated  in  another  way,  known 
as  the  amalgamation  process.  The  ores  are  mixed  with 
common  salt  and  roasted,  when  the  silver  is  obtained  in 
the  form  of  the  chloride.  This  is  then  reduced  to  silver 


METALLURGY  OF  SILVER.  603 

by  means  of  iron  and  water,  the  reaction  taking  place  as 
represented  in  the  following  equation  : 

2AgCl  +  Fe  =  FeCl2  +  2Ag. 

The  mixture  is  next  treated  with  mercury,  which  forms 
an  amalgam  with  silver,  while  the  other  metals  present 
do  not  combine  with  the  mercury.  The  amalgam  can 
be  separated  from  the  rest  of  the  mass  without  much 
difficulty,  and  when  heated  to  a  sufficiently  high  temper- 
ature the  mercury  distils  over,  leaving  the  silver. 

A  modification  of  the  amalgamation  process,  known  as 
the  American  process,  consists  in  grinding  the  ores  very 
fine,  mixing  them  with  sodium  chloride,  adding  roasted 
copper  pyrites,  which  consists  largely  of  copper  sul- 
phate, and  then  gradually  adding  mercury.  The  silver 
is  slowly  converted  into  silver  amalgam.  The  ex- 
planation of  the  process  is  this :  The  copper  sulphate 
reacts  with  the  sodium  chloride  to  form  cupric  chloride 
and  sodium  sulphate.  Cupric  chloride  reacts  upon  the 
silver  sulphide  as  represented  in  the  equation 

2CuCla  +  Ag2S  =  Cu2Cl2  +  2AgCl  +  S. 

The  cuprous  chloride  thus  formed  acts  upon  the  rest  of 
the  silver  sulphide,  forming  silver  chloride  and  cuprous 
sulphide  : 

Cu2Cl2  +  Ag2S  =  Cu2S  +  2AgCL 

The  silver  chloride  dissolves  in  sodium  chloride,  and  is 
then  reduced  and  converted  into  the  amalgam  by  the 
mercury. 

The  silver  in  the  market  is  not  pure.  For  chemical 
purposes  it  can  be  purified  by  dissolving  it  in  nitric 
acid,  precipitating  by  means  of  hydrochloric  acid,  filter- 
ing and  thoroughly  washing  the  chloride,  and  reducing 
this  either  by  melting  it  with  sodium  carbonate,  or  by 
pouring  a  little  dilute  hydrochloric  acid  upon  it,  and 
bringing  a  piece  of  zinc  in  contact  with  it.  In  the 
former  case  the  reaction  is 

2AgCl  +  Na2CO3  =  2Ag  +  C02  +  O  +  2NaCl ; 


604  INORGANIC  CHEMISTRY. 

in  the  latter  it  is 

Zn  +  2AgCl  =  ZnCl2  +  2Ag. 

Properties. — Silver  is  a  white  metal  with  a  high  lustre, 
of  specific  gravity  10.5.  It  is  not  acted  upon  by  the  air, 
oxygen,  or  water.  It  melts  at  a  lower  temperature  than 
copper  or  gold,  the  melting-point  being  about  1000°.  At 
the  temperature  of  the  oxyhydrogen  blowpipe  it  distils, 
and  in  the  experiments  of  Stas  on  the  atomic  weights  of 
chlorine  and  silver  the  metal  used  was  purified  in  this 
way.  It  is  harder  than  gold  and  softer  than  copper, 
and  its  "hardness  is  much  increased  by  the  addition  of  a 
little  copper.  It  combines  very  readily  with  sulphur, 
forming  black  silver  sulphide,  and  with  chlorine,  bro- 
mine, and  iodine.  The  blackening  of  silver  coins,  and 
other  objects  carried  about  the  person  is  caused  by  the 
presence  of  minute  quantities  of  sulphur  compounds  in 
the  perspiration ;  and  the  blackening  of  spoons  by  con- 
tact with  eggs  is  due  to  the  presence  of  sulphur  in  the 
albumen  of  the  eggs.  When  pure  silver  is  melted  in  the 
air  it  absorbs  about  twenty  times  its  volume  of  oxygen, 
and  this  is  given  off  when  the  metal  solidifies,  causing  in 
some  cases  a  sputtering  of  the  silver.  This  phenome- 
non is  observed  in  the  separation  of  silver  from  its  ores 
in  those  processes  in  which  it  is  necessary  to  melt  the 
metal.  It  is  known  as  "  spitting." 

At  the  ordinary  temperatures  silver  is  converted  into 
the  peroxide,  AgO,  by  ozone.  When  treated  with  hy- 
drochloric acid,  the  metal  becomes  covered  with  a  thin 
layer  of  the  chloride,  and  no  further  action  takes  place, 
but  it  is  dissolved  easily  by  concentrated  sulphuric  acid 
and  dilute  nitric  acid.  With  the  concentrated  acids  re- 
duction-products are  formed  as  with  copper.  Silver  is 
readily  dissolved  by  a  solution  of  potassium  cyanide  ; 
hence,  such  a  solution  is  used  in  removing  stains  caused 
by  silver  salts.  It  is  not  acted  upon  by  the  alkaline 
hydroxides  nor  by  potassium  nitrate  in  the  molten  con- 
dition, while  platinum  is.  Therefore,  silver  vessels  are 
used  when  it  is  desired  to  melt  these  substances  in  the 


SILVER  CHLORIDE.  605 

laboratory,   or  to  evaporate    their   solution,   as   in   the 
preparation  of  the  caustic  alkalies. 

Allotropic  Forms  of  Silver. — M.  Carey  Lea  has  dis- 
covered 'several  curious  allotropic  forms  of  silver,  the 
principal  of  which  are  briefly  described  by  him  as  fol- 
lows :  "A.  Soluble,  deep  red  in  solution,  mat  lilac,  blue, 
or  green  while  moist,  brilliant  bluish-green  metallic 
when  dry.  E.  Insoluble,  derived  from  A,  dark  red- 
dish-brown while  moist,  when  dry  somewhat  resembling 
A.  0.  Gold  silver,  dark  brown  while  wet,  when  dry 
exactly  resembling  metallic  gold  in  burnished  lumps. 
Of  this  form  there  is  a  variety  which  is  copper-colored. 
Insoluble  in  water;  appears  to  have  no  corresponding 
soluble  form." 

The  form  A  is  soluble  in  water,  and  the  solution  thus 
formed  has  a  deep  red  color.  The  different  varieties 
are  formed  by  the  action  of  reducing  agents  on  solutions 
of  silver  salts.  For  example,  the  red  soluble  form  is  ob- 
tained by  mixing  dilute  solutions  of  ferrous  citrate  and  a 
silver  salt.  All  the  allotropic  forms  of  silver  are  readily 
changed  to  the  ordinary  form. 

Mr.  Lea  further  says :  "  All  the  forms  of  allotropic 
silver  are  sensitive  to  light.  A  when  exposed  to  the 
sunlight  soon  becomes  brown.  The  bright  blue-green 
variety  of  B  is  changed  into  the  pure  gold-colored  variety 
of  C.  Other  forms  of  B  turn  brown  on  exposure  to 
light." 

The  red-yellow  variety  of  C  changes  to  bright  gold 
color.  "  Continued  exposure  seems  to  produce  little 
further  change  so  long  as  the  substance  is  dry.  But  if 
the  paper  on  which  the  silver  is  placed  is  kept  moist  by 
a  wet  pad,  with  three  or  four  days  of  good  sunshine,  the 
change  goes  on  until  the  silver  becomes  perfectly  white 
and  is  apparently  changed  to  normal  silver." 

Alloys  of  Silver. — For  practical  use,  as  in  making  coins 
and  silver-ware,  an  alloy  with  copper  is  used,  the  pure 
metal  being  too  soft.  The  alloy  usually  contains  from 
7^  to  10  per  cent  of  copper.  This  alloy  is  harder  than 
pure  silver,  and  is  capable  of  a  higher  polish.  Silver 


606  INORGANIC  CHEMISTRY. 

amalgam  is  an  alloy  of  silver  and  mercury,  which  is 
readily  formed  by  bringing  the  two  metals  together. 

Argentous  Chloride,  Ag2Cl  or  Ag4Cl3,  is  formed  by  treat- 
ing argentous  salts  with  hydrochloric  acid,  and,  possibly, 
to  some  extent  when  silver  chloride,  AgCl,  is  exposed  to 
the  light,  though  this  is  doubtful. 

Silver  Chloride,  Argentic  Chloride,  AgCl,  is  of  special 
importance  on  account  of  its  use  in  photography  and  in 
chemical  analysis.  It  occurs  in  nature  to  some  extent 
in  Mexico  and  in  the  United  States.  It  is  easily  formed 
as  a  white  precipitate  by  adding  hydrochloric  acid  to  a 
solution  of  a  silver  salt,  as,  for  example,  the  nitrate.  In 
consequence  of  its  insolubility  in  water  it  affords  a  con- 
venient means  of  detecting  silver  .and  chlorine.  If 
allowed  to  stand  in  the  light  it  changes  color,  becoming 
first  violet  and  finally  black.  This  change  in  color 
appears  to  be  due  entirely  to  the  reduction  of  the 
chloride  to  the  form  of  metallic  silver.  Concentrated 
hydrochloric  acid  dissolves  it  somewhat,  and  from  this 
solution  it  crystallizes  in  octahedrons.  An  aqueous  solu- 
tion of  ammonia  dissolves  it  very  easily,  in  consequence 
of  the  formation  of  a  compound  of  the  chloride  with  am- 
monia analogous  to  those  formed  by  copper  salts.  The 
composition  of  the  compound  in  the  solution  is,  how- 
ever, not  known.  Concentrated  solutions  of  potassium, 
sodium,  and  ammonium  chlorides  dissolve  silver  chlo- 
ride, forming  double  chlorides ;  and  potassium  cyanide 
also  forms  an  easily  soluble  double  salt  with  it.  The 
dry  compound  absorbs  ammonia  gas,  forming  a  com- 
pound of  the  formula  2AgC1.3NH3,  which  readily  gives 
up  the  ammonia  when  gently  heated. 

Silver  Bromide,  AgBr,  and  Silver  Iodide,  Agl,  are  very 
similar  to  the  chloride.  Both  occur  in  nature,  and  both 
are  precipitated  from  solutions  of  silver  salts  by  adding 
the  corresponding  hydrogen  acids.  The  bromide  is  less 
easily  soluble  in  ammonia  than  the  chloride,  and  the 
iodide  is  almost  insoluble  in  it.  The  bromide  is  formed 
by  treating  the  chloride  at  the  ordinary  temperature  with 
hydrobromic  acid;  and  the  iodide  is  formed  from  the 
chloride  and  from  the  bromide  by  treating  these  with 


8ALT8  OF  SILVER  IN  PHOTOGRAPHY.  607 

hydriodic  acid  at  ordinary  temperatures.  At  higher 
temperatures,  however,  both  the  bromide  and  iodide  are 
converted  into  the  chloride  by  hydrochloric  acid.  Silver 
dissolves  in  concentrated  hydriodic  acid,  and  from  the 
solution  a  salt  of  the  formula  Agl  -\-  HI  or  HAgIa  is 
formed.  It  seems  probable  that  this  is  a  derivative  of 
the  acid  H2I2,  from  which  the  double  salt  KI.AgI  is  also 
derived,  as  indicated  in  the  formula  KAgI2.  Silver 
bromide  at  low  temperatures  is  white,  but  easily  changes 
to  yellow,  and  by  exposure  it  becomes  darker,  but  not  as 
readily  as  the  chloride.  The  iodide  is  yellow,  and  under- 
goes change  in  the  light  only  very  slowly.  The  chloride 
and  iodide  exist  in  several  modifications,  which  differ 
from  one  another  in  their  conduct  towards  light,  and  in 
their  solubility.  Probably  the  differences  are  due  to 
different  complexity  of  the  molecules.  Modifications 
corresponding  to  the  formulas  AgCl,  Agl,  Ag2Cl2,  Ag2I2, 
Ag3Cl3,  Ag3I3,  etc.,  are  quite  conceivable.  The  careful 
study  of  the  effects  of  light  upon  the  different  modifica- 
tions seems  to  promise  interesting  results,  which  may 
make  it  possible  to  judge  as  to  the  relative  complexity 
of  the  molecules. 

Application  of  the  Chloride,  Bromide,  and  Iodide  of 
Silver  in  the  Art  of  Photography. — The  art  of  photography 
is  based  upon  the  changes  which  certain  compounds, 
especially  salts  of  silver,  undergo  when  exposed  to  the 
light.  Silver  iodide  is  best  adapted  to  most  purposes. 
The  salt  is  so  changed  by  the  light  that  when  treated 
with  certain  compounds,  such  as  ferrous  sulphate,  pyro- 
gallic  acid,  etc.,  called  "  developers,"  a  deposit  of  finely 
divided  silver  is  formed  upon  the  plate  in  those  places 
affected  by  the  light.  A  plate  of  glass  or  a  sheet  of 
properly  prepared  paper  is  covered  in  the  dark  with  a 
thin  layer  of  a  salt  of  silver.  The  plate  is  then  exposed 
in  the  camera  to  the  action  of  the  light  which  is  reflected 
from  the  object  to  be  photographed.  According  to  the 
intensity  of  the  light  given  off  from  the  various  parts  of 
the  object,  the  change  of  the  silver  salt  takes  place  to  a 
greater  or  less  extent,  and  thus  a  perfect  image  of  the 
object  is  impressed  upon  the  plate.  But  after  the  action 


608  INORGANIC  CHEMISTRY. 

of  the  "  developer"  is  complete  there  is  still  upon  the 
plate  unchanged  silver  salt,  and  if  this  were  now  exposed 
to  the  light  it  would  undergo  change  and  the  image 
would  be  obliterated.  To  remove  this  salt  the  plate  is 
washed  with  a  solution  of  sodium  thiosulphate,  Na2S2O3 
(hyposulphite),  which  dissolves  the  salt  in  consequence 
of  the  formation  of  a  double  salt  of  the  formula 
2Na2S2O3.Ag2S2O3,  which  is  readily  soluble  in  water. 

Silver  Triazoate,  AgN3. — This  is  derived  from  triazoic 
acid  (which  see).  It  is  formed  by  adding  a  solution  of  the 
acid  to  a  solution  of  a  silver  salt.  It  is  extremely  explo- 
sive and  should  be  dealt  with  very  cautiously.  Serious 
accidents  have  been  caused  by  it.  In  appearance  it 
resembles  silver  chloride,  but  it  does  not  darken  when 
exposed  to  the  light. 

Silver  Oxide,  Ag2O. — The  principal  compound  of  silver 
and  oxygen  is  that  which  has  the  composition  Ag2O,  and 
in  which  the  silver  is  univalent,  as  it  is  in  its  compounds 
with  chlorine,  bromine,  and  iodine.  It  is  formed  when 
a  soluble  hydroxide  is  added  to  a  solution  of  a  silver 
salt,  and  also  by  the  action  of  concentrated  solutions  of 
the  caustic  alkalies  on  silver  chloride.  It  is  easily  de- 
composed by  heat  and  by  reducing  agents. 

Other  Oxides  of  Silver. — Besides  the  ordinary  oxide, 
silver  forms  a  sub-oxide,  Ag4O,  corresponding  to  the  sub- 
oxide  of  copper,  Cu4O,  and  a  peroxide  of  the  formula 
AgO  (or  Ag4O3),  which  is  perhaps  analogous  to  cupric 
oxide. 

Sulphides  of  Silver. — As  has  been  stated,  silver  occurs 
in  nature  mostly  in  combination  with  sulphur  as  silver 
glance,  Ag2S,  which  is  in  many  minerals  in  combination 
with  other  sulphides.  Examples  of  such  double  sul- 
phides are  the  minerals  stromeyerite,  Cu2S.Ag2S,  and 
pyrargyrite,  3Ag2S.Sb2S8. 

Silver  Nitrate,  Argentic  Nitrate,  AgNO3. — This  salt  is 
formed  by  dissolving  silver,  or  silver  oxide,  in  nitric  acid, 
evaporating  to  dryness,  and  heating  until  the  salt  is 
melted.  It  crystallizes  in  colorless  rhombic  plates.  It 
is  not  changed  in  the  light  unless  it  comes  in  contact 


COMPOUNDS  OF  SILVER.  609 

with  organic  substances,  when  it  is  reduced  and  metallic 
silver  deposited.  Hence  the  solution  produces  black 
spots  on  the  fingers  and  clothing.  As  it  melts  easily,  it 
is  generally  cast  in  small  cylindrical  moulds,  and  is  found 
in  the  market  in  the  form  of  thin  sticks,  and  is  known  as 
lunar  caustic.  It  disintegrates  flesh,  and  is  used  in  sur- 
gery as  a  caustic  to  remove  superfluous  growths.  Owing 
to  the  formation  of  a  dark  deposit  when  the  salt  is  ex- 
posed to  the  light,  it  is  used  as  a  constituent  of  indelible 
inks.  The  dry  nitrate  absorbs  ammonia  and  forms  the 
compound  AgNO3  +  3NH3 ;  in  concentrated  solution  the 
compound  AgNOs  +  2NH3  is  formed. 

Silver  Cyanide,  AgCN,  is  formed  as  a  caseous  pre- 
cipitate when  a  solution  of  hydrocyanic  acid  is  added  to 
a  solution  of  silver  nitrate.  It  does  not  change  color  in 
the  light,  is  soluble  in  ammonia,  but  hot  in  nitric  acid. 
It  readily  forms  double  cyanides  with  the  cyanides  of 
other  metals.  Of  these,  the  salt  with  potassium  cyanide, 
KAg(CN)2  or  KCN.AgCN,  may  be  mentioned. 

Silver  Sulphocyanate,  AgSCN,  is  very  similar  to  the 
cyanide,  and  is  formed  when  solutions  of  silver  nitrate 
and  potassium  or  ammonium  sulphocyanate  are  brought 
together.  It  is  soluble  in  an  excess  of  the  soluble  cy- 
anides, double  salts  similar  to  the  double  cyanides  being 
formed. 

Borates  of  Silver. — When  a  cold  concentrated  solution 
of  sodium  metaborate,  NaBO2,  is  mixed  with  a  similar 
solution  of  silver  nitrate  a  precipitate  of  silver  meta- 
borate, AgBO2,  containing  some  silver  oxide  is  formed. 
When  dilute  solutions  of  the  two  compounds  are  mixed 
a  precipitate  of  silver  oxide  is  formed  ;  so,  also,  silver 
metaborate  is  decomposed  by  water  into  boric  acid  and 
silver  oxide,  and  when  the  solution  in  which  the  pre- 
cipitate is  suspended  is  boiled  the  same  change  takes 
place.  Further,  when  cold  concentrated  solutions  of 
silver  nitrate  and  borax  are  mixed,  silver  octoborate, 
Ag6B8O15,  is  precipitated,  and  this  is  mixed  with  some 
silver  oxide.  When  the  solution  is  boiled,  the  silver  salt 
is  decomposed  into  boric  acid  and  silver  oxide.  When 


610  INORGANIC  CHEMISTRY. 

the  solutions  of  borax  and  silver  nitrate  are  mixed  hot, 
the  precipitate  is  the  metaborate  of  silver. 

Reactions  which  are  of  Special  Value  in  Chemical 
Analysis. — Hydrochloric  acid  precipitates  insoluble  silver 
chloride  from  solutions  of  silver  salts,  as  silver  nitrate. 

Soluble  hydroxides  precipitate  silver  oxide,  not  the  hy- 
droxide. Ammonia  redissolves  the  precipitate  in  conse- 
quence of  the  formation  of  a  compound  of  the  oxide  with 
ammonia  of  the  composition  Ag2O.2NH3.  In  dry  con- 
dition this  salt  is  very  explosive,  and  is  known  as  ful- 
minating silver. 

Soluble  carbonates  precipitate  the  carbonate,  Ag2CO3, 
which  has  a  yellowish-white  color. 

Ammonium  carbonate  redissolves  the  precipitate  formed 
by  it. 

Sodium  phosphate,  HNaJPO4,  gives  a  precipitate  of  the 
normal  salt  Ag3PO4 ,  which  is  yellow. 

Potassium  f err ocyanide  precipitates  white  silver  ferro- 
cyanide,  Ag4Fe(CN)6. 

Potassium  ferricyanide,  K3Fe(CN)6 ,  gives  the  corre- 
sponding silver  salt,  which  is  reddish  brown. 

Potassium  chromate  or  potassium  dichromate  (which  see) 
gives  a  brownish-red  precipitate  of  silver  chromate. 

GOLD,  Au  (At.  Wt.  195.74). 

General. — Gold  forms  two  series  of  compounds,  in  one 
of  which  it  is  univalent  and  in  the  other  trivalent.  In 
this  respect  it  differs  from  the  other  members  of  the 
group.  Examples  of  the  compounds  belonging  to  the 
two  series  are  represented  by  the  following  formulas : 

AuCl  AuCl3 

AuBr  AuBr3 

Au2O  Au2O3 

Those  of  the  first  series  are  called  aurous  compoundst 
those  of  the  second  series  auric  compounds.  The  basic 
character  of  gold  is  very  weak,  so  that  salts  of  the  ordi- 
nary acids,  as  sulphuric,  nitric,  carbonic,  etc.,  are  not 


GOLD.  611 

known.  On  the  other  hand,  its  higher  oxide  and  hy- 
droxide, Au(OH)3,  have  acid  properties,  and  form  salts 
similar  in  composition  to  the  meta-aluminates  MAlOa, 
and  the  metaborates  MBO2.  These  are  the  aurates,  of 
which  potassium  aurate,  KAuO2,  is  an  example.  So,  also, 
the  chloride  combines  readily  with  the  chlorides  of  potas- 
sium and  sodium,  forming  the  chlor-aurates,  KAuCl4,  and 
NaAuCl4,  which  are  perfectly  analogous  to  the  aurates. 
Further,  the  chloride  and  bromide  combine  respectively 
with  hydrochloric  and  hydrobromic  acids,  forming  the 
crystallized  compounds  HAuCl4  +  4H2O  and  HAuBr4  + 
5H2O,  which  are  plainly  the  acids  from  which  the  chlor- 
aurates  and  the  brom-aurates  are  derived. 

Besides  the  compounds  of  gold  in  which  the  element 
is  univalent  and  those  in  which  it  is  tnvalent,  the  chlo- 
ride AuCl3,  and  the  bromide,  AuBr2,  have  been  described, 
but  their  existence  is  doubtful. 

Forms  in  which  Gold  occurs  in  Nature. — Gold  is  gen- 
erally found  in  nature  in  the  native  condition — a  fact 
which  is  undoubtedly  due  to  the  chemical  inactivity  of 
the  element.  That  which  is  found  in  nature  is  never 
pure,  but  contains  silver,  and  also,  in  different  localities, 
iron,  copper,  and  other  metals.  It  is  also  found  to  some 
extent  in  combination  with  tellurium  in  the  compounds 
AuTe2  and  (AuAg)2Te3.  Native  gold  is  frequently  found 
enclosed  in  quartz,  or  more  commonly  in  quartz  sand. 
The  principal  localities  in  which  it  is  found  are  California 
and  some  of  the  other  Western  United  States,  and 
Australia,  Hungary,  Siberia,  and  Africa. 

Metallurgy  of  Gold. — From  the  chemical  point  of  view 
the  metallurgy  of  gold  is  in  general  very  simple.  There 
are  two  kinds  of  gold  mining — called  placer  mining  and 
vein  mining.  In  the  former  the  earth  and  sand  which 
contain  gold  are  washed  with  water,  which  carries  away 
the  lighter  particles,  and  leaves  the  gold  mixed  with 
other  heavy  materials.  This  mixture  is  then  treated 
with  mercury,  which  forms  an  amalgam  with  the  gold,  as 
it  does  with  silver,  and  when  this  is  placed  in  a  prop- 
erly constructed  retort  and  heated,  the  mercury  passes 
over  and  leaves  the  gold  behind.  If  silver  is  present,  as 


612  INORGANIC  CHEMISTRY. 

is  frequently  the  case,  this  is  separated  with  the  gold, 
In  vein  mining  the  gold  ores  are  taken  out  of  veins  in  the 
earth,  and  the  gold  separated  by  grinding  the  ores  and 
treating  them  with  mercury,  as  in  the  last  stage  of  placer 
mining.  Hydraulic  mining  is  a  modification  of  ordinary 
placer  mining.  It  consists  in  forcing  water  under  pres- 
sure against  the  sides  of  hills  and  mountains  in  which 
gold  occurs  loosely  mixed  with  the  earth.  The  earth  is 
thus  carried  away  and  the  heavier  gold  is  deposited  in 
sluices. 

Some  ores,  those  especially  which  contain  tellurium, 
cannot  be  satisfactorily  treated  by  the  amalgamation 
process,  and  a  method  involving  the  use  of  potassium 
cyanide  has  been  devised  for  them.  In  a  solution  of  this 
salt  gold  dissolves,  and  from  this  solution  it  can  be  sep- 
arated in  various  ways.  This  method  has  come  into  ex- 
tensive use  of  late  years. 

Another  process  that  is  extensively  used  in  the  treat- 
ment of  ores  that  do  not  give  their  gold  to  mercury  is 
known  as  the  chlorination  process.  This  consists  in  treat- 
ing the  finely  ground  ore  with  chlorine  made  from 
bleaching  powder  and  sulphuric  acid,  and  then  pre- 
cipitating the  gold  from  the  solution  of  the  chloride  by 
means  of  hydrogen  sulphide.  From  the  sulphide  the 
metallic  gold  can  be  easily  obtained. 

The  gold  obtained  by  any  of  the  above  methods  is  not 
pure.  It  can  be  separated  from  silver  by  dissolving  it 
in  aqua  regia,  evaporating  so  as  to  drive  off  the  nitric 
acid,  then  diluting,  and  treating  with  a  reducing  agent, 
when  metallic  gold  is  precipitated.  Thus  when  ferrous 
sulphate  is  used  the  following  reaction  takes  place : 

3FeSO4  +  AuCl3  =  Fe2(SO4)3  +  FeCl3  +  Au. 

Another  method  of  separating  silver  from  an  alloy 
with  gold  consists  in  treating  the  metal  with  nitric  acid 
or  with  boiling  concentrated  sulphuric  acid,  which  dis- 
solves the  silver  and  leaves  the  gold.  This  process  is 
not  satisfactory,  however,  unless  the  amount  of  gold  in 
the  alloy  is  less  than  25  per  cent.  If  the  proportion  of 
gold  is  greater  than  this,  the  alloy  is  melted  with  silver 


PROPERTIES  OF  GOLD— ALLOYS— CHLORIDES.     613 

enough  to  bring  the  percentage  of  gold  down  to  that 
mentioned.  This  is  known  as  " quartation" 

Properties. — Gold  is  a  yellow  metal  with  a  high  lustre. 
It  is  quite  soft,  and  extremely  malleable,  so  that  it  is 
possible  to  make  from  it  sheets  the  thickness  of  which  is 
not  more  than  0.000002  millimeter.  Thin  sheets  are 
translucent,  and  the  transmitted  light  appears  green. 
Its  specific  gravity  is  19.3  ;  its  melting-point  higher  than 
that  of  copper,  being  about  1200°.  It  crystallizes  in  the 
regular  system.  Gold  combines  directly  with  chlorine, 
but  not  with  oxygen.  The  three  acids,  hydrochloric, 
nitric,  and  sulphuric,  do  not  act  upon  it ;  but  aqua 
regia  dissolves  it,  forming  auric  chloride,  AuCl3,  in  con- 
sequence of  the  evolution  of  nascent  chlorine.  Molten 
caustic  alkalies  and  their  nitrates  act  upon  it,  probably 
in  consequence  of  the  tendency  to  form  aurates. 

Alloys  of  Gold. — The  principal  alloy  of  gold  is  that 
which  contains  copper.  The  standard  gold  coin  of  the 
"United  States  contains  nine  parts  of  gold  to  one  of  cop- 
per. The  composition  of  gold  used  for  jewelry  is  usually 
stated  in  terms  of  carats.  Pure  gold  is  24-carat  gold ; 
20-carat  gold  contains  20  parts  of  gold  and  4  parts  of 
copper ;  18-carat  gold  contains  18  parts  of  gold  and  6 
parts  of  copper,  etc.  Copper  gives  gold  a  reddish  color, 
and  makes  it  harder  and  more  easily  fusible.  Gold  is 
also  alloyed  with  silver ;  and  the  alloy  with  mercury, 
known  as  gold-amalgam,  is  extensively  used  in  the  pro- 
cesses for  extracting  gold  from  its  ores. 

Chlorides  of  G-old. — When  gold  is  dissolved  in  aqua 
regia  it  is  converted  into  auric  chloride,  AuGl3 ;  and  if  this 
solution  is  evaporated  a  part  of  the  chloride  is  decom- 
posed into  aurous  chloride,  AuCl,  and  chlorine.  When  gold 
is  treated  with  dry  chlorine  it  yields  a  mixture  of  auric 
chloride  and  metallic  gold.  This  was  formerly  held  to 
be  a  chloride  of  the  formula  AuCl2,  but  the  most  careful 
investigations  on  the  subject  have  shown  that  this  does 
not  exist.  Auric  chloride  can  be  obtained  in  crystal- 
lized form,  the  crystals  having  the  composition  AuCl8  -[- 
2H2O.  When  anhydrous  auric  chloride  is  heated  to 
185°,  it  loses  chlorine  and  is  converted  into  aurous  chlo- 


614  INORGANIC  CHEMISTRY. 

ride,  AuCl.  This  yields  auric  chloride  and  gold  when 
treated  with  water.  When  treated  with  a  solution  of 
stannous  chloride  a  solution  of  auric  chloride  gives  a 
purple-colored  precipitate,  known  as  the  purple  of  Cas- 
sius,  which  appears  to  consist  of  finely-divided  gold. 

Chlor-auric  Acid  and  its  Salts.  —  When  a  solution  of 
gold  in  aqua  regia  containing  a  large  excess  of  hydro- 
chloric acid  is  evaporated  a  crystallized  product  of  the 
formula  HAuCl4  +  4H2O,  or  AuCl3.HCl  +  4H3O,  is  ob- 
tained. This  is  chlor-auric  acid.  It  must  be  regarded 
as  belonging  to  the  same  class  as  fluosilicic  acid  and 
the  chloro-acids,  from  which  the  double  chlorides  of 
magnesium,  aluminium,  copper,  etc.,  are  derived.  Ac- 
cordingly its  constitution  is  expressed  by  the  formula 

/ci 

Au^-Cl       ,  being  similar  to  that  of  the  acid  from  which 


/Cl 
potassium  chlor-aluminate  is  derived,  A1—C1        .     The 

\(C1,)H 

potassium  salt,  KAuCl4,  is  obtained  by  mixing  together 
solutions  of  auric  and  potassium  chlorides. 

Cyan-auric  Acid,  HAu(CN)4,  is  perfectly  analogous  to 
chlor-auric  acid.  It  is  formed  by  treating  the  potassium 
salt,  KAu(CN)4,  with  silver  nitrate,  which  gives  the  silver 
salt,  and  then  decomposing  this  with  hydrochloric  acid. 
The  potassium  salt  is  obtained  by  mixing  solutions  of 
auric  chloride  and  potassium  cyanide.  The  salts  of  a 
cyan-aurous  acid,  HAu(CN)2,  are  also  known. 

Auric  Hydroxide,  Au(OH)3.  —  This  compound  is  formed 
by  treating  a  solution  of  auric  chloride  with  an  excess  of 
magnesia  or  with  sodium  hydroxide,  and  afterwards  with 
sodium  sulphate.  It  is  a  yellow  or  brown  powder. 
When  exposed  to  the  light  it  is  decomposed  with  evolu- 
tion of  oxygen.  When  heated  to  100°  it  yields  auric 
oxide,  Au2O3,  and  when  this  is  heated  to  a  higher  tempera- 
ture it  loses  all  its  oxygen.  Aurous  oxide,  Au2O,  is  formed 
by  treating  aurous  chloride  with  caustic  potash.  It  is 
easily  decomposed  by  heat  into  gold  and  oxygen. 

Auric  hydroxide  dissolves  in  the  soluble  hydroxides 
just  as  aluminium  hydroxide  does,  and  from  the  solutions 


AURIC  HYDROXIDE—  GOLD  SULPHIDE.  615 

salts  known  as  the  aurates  are  obtained.  In  composition 
these  are  analogous  to  the  meta-aluminates.  Potassium 
aurate,  for  example,  has  the  composition  KAuO2.  The 
analogy  between  some  of  the  compounds  of  aluminium 
and  those  of  gold  is  shown  in  the  following  table  : 

A1A  Au203 

A1(OH)3  Au(OH)3 


A1C1,  AuCl3 

/Cl  /Cl 

Al(-Cl  Au^-Cl 
\(C12)K  \(C13)K 

Gold  Sulphide,  Au2S2.  —  This  compound  is  precipitated 
together  with  sulphur  from  cold  solutions  of  gold  salts 
by  means  of  hydrogen  sulphide,  and  forms  a  brownish 
black  mass.  It  forms  soluble  compounds  with  the  sul- 
phides of  the  alkali  metals. 

When  hydrogen  sulphide  is  passed  into  hot  solutions 
of  gold  salts  aurous  sulphide,  Au2S,  is  thrown  down  as 
a  steel-gray  substance.  This  is  soluble  in  pure  water 
and  is  reprecipitated  by  hydrochloric  acid, 


CHAPTER  XXIX. 

ELEMENTS  OF  FAMILY  II,  GROUP  B: 
ZINC— CADMIUM— MERCURY. 

General. — There  is  a  very  strong  resemblance  between 
the  first  two  elements  of  this  group  and  magnesium, 
while  mercury,  in  a  general  way,  resembles  the  first  two 
members  of  the  copper  group.  Just  as  gold  in  the  cop- 
per group  furnishes  a  greater  variety  of  compounds  than 
the  first  two  members  of  that  group,  so  mercury  fur- 
nishes a  greater  variety  of  compounds  than  the  other 
members  of  the  group  to  which  it  belongs.  Zinc  and 
cadmium,  like  magnesium,  give  only  one  class  of  com- 
pounds and  in  these  they  are  bivalent.  The  general  for- 
mulas of  some  of  the  principal  ones  are : 

MC12,    M(OH)2,    MO,    MSO4,    MCO8,    MS. 

Mercury,  on  the  other  hand,  furnishes  two  series  of  com- 
pounds, known  as  the  mercurous  and  mercuric  compounds, 
which  correspond  closely  to  the  two  series  of  copper 
salts.  The  power  to  form  compounds  belonging  to  both 
series  is  more  strongly  developed  in  mercury  than  in. 
copper.  Examples  of  the  two  classes  are  represented  in 
the  following  formulas  : 

Mercurous  Compounds.  Mercuric  Compounds. 

HgCl  HgCl2 

Hgl  HgI2 

Hg20  HgO 

HgN03  Hg(N03)2,    etc. 

Just  as  the  first  member  of  Group  A,  Family  II,  gluci- 
num,  shows  a  somewhat  acidic  character  in  its  hydrox- 
ide, while  the  other  members  of  that  group  do  not;  so 
also  the  first  member  of  Group  Bj  Family  II,  zinc,  is 

(616) 


ZING.  617 

acidic,  while  the  other  members  of  the  group  are  not. 
Glucinum  hydroxide  and  zinc  hydroxide  dissolve  in 
caustic  alkalies,  forming  glucinates  and  zincates ;  while 
the  hydroxides  of  all  the  other  members  of  the  two 
groups  of  this  family  are  insoluble  in  caustic  alkalies. 

ZINC,  Zn  (At.  Wt.  64.91). 

General. — Zinc,  in  almost  all  its  compounds,  exhibits 
a  close  resemblance  to  magnesium.  It  always  acts  as  a 
bivalent  element. 

Forms  in  which,  it  occurs  in  Nature. — Zinc  occurs  in 
nature  in  combination  principally  as  the  carbonate,  or 
smithsonite,  ZnCO3 ;  as  the  sulphide,  or  sphalerite,  ZnS  ; 
and  as  the  silicate,  Zn2SiO4.  Among  other  compounds 
of  zinc  found  in  nature  are  gahnite,  Zn(AlO2)2,  and  frank- 
linite,  which  contains  the  compound  Zn(FeO2)2  with 
Fe(Fe02)2. 

Metallurgy. — The  metallurgy  of  zinc  is  much  simpler 
than  that  of  magnesium,  for  the  reason  that  the  ores  are 
easily  converted  into  the  oxide  by  roasting,  and  the  oxide 
is  easily  reduced  by  heating  it  with  charcoaL  Owing  to 
the  volatility  of  the  metal  the  vessels  in  which  the  reduc- 
tion is  effected  must  be  so  constructed  as  to  facilitate  the 
condensation  of  the  vapors.  The  vessels  used  are  either 
earthenware  muffles  or  tubes,  open  at  one  end  and  con- 
nected with  iron  receivers.  At  first  the  zinc  vapor  is 
condensed  in  the  form  of  a  fine  dust,  as  in  the  case  of 
sulphur.  This  forms  the  commercial  product  called 
zinc  dust.  It  always  contains  zinc  oxide.  Afterwards 
the  zinc  condenses  to  the  form  of  a  liquid,  and  this  is 
cast  in  plates.  The  zinc  thus  obtained  is  not  pure,  but 
contains  lead  and  iron,  and  sometimes  arsenic  and  cad- 
mium. It  is  called  spelter.  By  repeated  distillation  it 
can  be  obtained  pure.  When  distilled  under  diminished 
pressure,  it  is  deposited  in  beautiful  lustrous  crystals, 
the  forms  of  which  are  extremely  complicated. 

Properties. — Zinc  has  a  bluish-white  color  and  a  high 
lustre.  The  crystals  above  referred  to,  which  are  per- 
fectly pure  zinc,  have  a  brilliant  lustre,  and  do  not 


618  INORGANIC  CHEMISTRY. 

change  in  the  air.  At  different  temperatures  zinc  has 
markedly  different  properties.  At  ordinary  temperatures 
it  is  quite  brittle ;  at  100°-150°  it  can  be  rolled  out  in 
sheets,  but  above  200°  it  becomes  brittle  again.  It  melts 
at  433°,  and  boils  at  1040°.  When  heated  in  the  air  it 
takes  fire,  and  burns  with  a  bluish  flame,  forming  zinc 
oxide.  This  can  be  shown  by  means  of  the  oxyhydro- 
gen  blowpipe.  In  dry  air  it  does  not  change.  Ordinary 
zinc  dissolves  in  all  the  common  acids,  usually  with  evo- 
lution of  hydrogen.  In  the  case  of  nitric  acid,  however, 
the  hydrogen  acts  upon  the  acid,  reducing  it  to  ammonia. 
The  purer  the  zinc  the  less  readily  is  it  acted  upon  by 
sulphuric  acid,  and  the  pure  crystals  above  referred 
to  are  scarcely  acted  upon  at  all  by  this  acid.  Zinc 
also  dissolves  in  the  caustic  alkalies,  forming  zincates. 
Pure  zinc  can  be  made  to  act  upon  sulphuric  acid  by 
adding  a  few  drops  of  platinum  chloride. 

Applications. — Zinc  is  extensively  used  as  sheet-zinc, 
in  making  galvanic  batteries,  for  galvanizing  iron,  etc. 
Zinc  dust  is  a  very  efficient  reducing  agent,  either  in  al- 
kaline or  in  acid  solution.  With  caustic  alkalies — as,  for 
example,  with  potassium  hydroxide — it  gives  hydrogen ' 
and  a  zincate : 

Zn  +  2KOH  =  Zn(OK)2  +  H8. 

With  sulphuric  acid  also  it  gives  hydrogen  readily. 
Zinc  is  used  in  the  preparation  of  important  alloys. 

Alloys. — Iron  covered  with  a  layer  of  zinc  is  known 
as  galvanized  iron.  As  has  been  mentioned,  zinc  is  a 
constituent  of  brass.  It  combines  readily  with  mercury 
to  form  zinc  amalgam,  and  this  fact  is  taken  advantage 
of  for  the  purpose  of  preserving  the  zinc  plates  in  gal- 
vanic batteries.  Zinc  plates  covered  with  a  layer  of 
the  amalgam  are  acted  upon  much  more  slowly  than 
zinc  itself.  The  amalgamation  is  effected  by  cleaning 
the  zinc,  dipping  it  in  dilute  sulphuric  acid,  and  rubbing 
mercury  over  the  surface  with  a  brush  or  a  piece  of 
cloth. 

Zinc  Chloride,  ZnCl2. — This  is  prepared  by  treating 
zinc  with  chlorine,  or  by  dissolving  zinc  in  hydrochloric 


ZINC  CHLORIDE—  ZINC  HYDROXIDE.  619 

v 

acid,  evaporating  to  dryness,  and  distilling  the  residue.  It 
is  a  white  deliquescent  mass.  From  a  very  concentrated 
solution  in  hydrochloric  acid  it  is  obtained  in  crystals 
of  the  composition  ZnCl2  -f-  H2O.  When  the  solution  is 
evaporated  there  is  always  some  decomposition  into  basic 
chlorides,  the  hydroxide,  and  oxide.  The  basic  chloride 
is  formed  thus  : 


HHO  =  Zn<         +  HC1  ; 

the  hydroxide  thus  : 

HHO  =  Zn<  HC1 


and  by  higher  heating  the  hydroxide  yields  the  oxide 
and  water  : 

Zn<OH  =  Zn°  +  HA 

The  chloride  has  a  marked  affinity  for  water,  and  is  used 
in  the  laboratory,  as  sulphuric  acid  and  phosphorus 
pentoxide  are,  for  the  purpose  of  extracting  the  elements 
of  water  from  compounds.  It  has  a  caustic  action,  and 
is  used  in  surgery  on  this  account.  Further,  it  acts  as  a 
disinfectant,  and  its  solution  is  used  for  the  purpose  of 
preserving  wood,  particularly  railroad  sleepers,  from  de- 
cay. 

The  chloride  readily  forms  double  chlorides  like  those 
formed  by  magnesium  chloride.  Examples  of  these  are 
the  compounds  of  the  formulas  ZnCl2.2KCl  or  K2ZnCl4, 
ZnCl2.2NaCl  or  Na2ZnCl4,  etc.  The  double  chloride 
with  ammonium  chloride,  (NH4)2ZnCl4,  is  formed  by 
mixing  a  solution  of  zinc  in  hydrochloric  acid  with  a 
solution  of  ammonium  chloride.  This  is  used  in  solder- 
ing, as  it  cleans  the  surface  of  the  metal,  in  consequence 
of  the  action  of  the  zinc  chloride  on  the  oxides.  It  also 
absorbs  ammonia,  forming  compounds  analogous  to  those 
formed  by  cupric  and  cuprous  chlorides. 

Zinc  Hydroxide,  Zn(OH)2,  is  precipitated  as  a  white 
amorphous  powder  when  a  soluble  hydroxide  is  added 
to  a  solution  of  a  zinc  salt.  It  is  redissolved  in  an  ex- 


620  INORGANIC  CHEMISTRY. 

cess  of  the  reagent,  and  the  zincate  thus  formed  is  de- 
composed on  boiling,  the  hydroxide  being  reprecipitated. 

Zinc  Oxide,  ZnO,  is  formed  in  very  finely  divided  con- 
dition  by  burning  zinc  in  the  air.  The  product  is  known 
as  Flores  zinci,  and  is  sometimes  called  philosopher's 
wool.  It  is  also  formed  by  heating  the  carbonate  or  ni- 
trate, and  is  found  in  nature  mixed  with  or  in  combina- 
tion with  oxide  of  manganese,  Mn3O4.  It  is  prepared  on 
the  large  scale  both  by  burning  zinc  and  by  heating  the 
basic  carbonate,  which  is  formed  by  adding  sodium  car- 
bonate to  a  solution  of  zinc  sulphate.  It  is  a  white  pow- 
der, which  turns  yellow  when  heated.  Its  chief  use  is  as 
a  constituent  of  paint  under  the  name  of  zinc  ivhite. 

Zinc  Sulphide,  ZnS. — This  compound  occurs  in  nature, 
and  is  known  as  zinc  blende.  The  mineral  always  con- 
tains a  sulphide  of  iron,  and  also  a  small  quantity  of 
cadmium  sulphide.  When  hydrogen  sulphide  is  passed 
into  a  solution  of  a  zinc  salt  only  a  part  of  the  zinc  is 
thrown  down  as  the  sulphide,  if  the  salt  used  is  one  of  a 
strong  acid,  like  sulphuric,  nitric,  or  hydrochloric  acid. 
The  reason  of  this  is  that  the  sulphide  is  soluble  in  these 
acids,  even  when  they  are  very  dilute.  In  the  reaction 
the  acid  is  set  free,  and  although  some  sulphide  is  thrown 
down,  the  action  soon  stops : 

ZnSO4  +  H2S  =  ZnS  +  H2SO4. 

If  the  acetate  of  zinc  is  used  the  precipitation  is  com- 
plete, because  dilute  acetic  acid  does  not  dissolve  zinc 
sulphide.  If  sodium  or  potassium  acetate  is  added  to  a 
solution  of  a  neutral  salt  of  zinc,  hydrogen  sulphide  pre- 
cipitates all  the  zinc,  for  the  reason  that  the  acid  which  is 
first  set  free  acts  upon  the  acetate  and  is  itself  neu- 
tralized, while  acetic  acid  is  then  set  free.  Thus,  when 
hydrogen  sulphide  acts  upon  a  solution  of  zinc  sulphate 
containing  sodium  acetate  the  action  involves  two  steps, 
as  represented  in  the  two  equations  : 

ZnSO4  +  H2S  =  ZnS  +  H2SO4 ; 
2NaC,H302  +  H2S04  =  Na2SO4  +  2C2H4Oa. 


COMPOUNDS  OF  ZINC.  621 

As  fast  as  the  sulphuric  acid  is  formed  it  acts  upon  the 
aeetate,  and  is  thus  prevented  from  dissolving  the  sul- 
phide. 

The  sulphide  is,  further,  completely  precipitated  by 
soluble  sulphides,  as  potassium  and  ammonium  sul- 
phides. Obtained  by  precipitation,  zinc  sulphide  is  a 
white  amorphous  substance. 

Zinc  Sulphate,  ZnSO4.  —  This  salt  is  readily  formed  by 
oxidation  of  the  sulphide,  and  is  hence  found  in  nature 
accompanying  the  sulphide.  It  is  manufactured  by  care- 
fully roasting  zinc  blende,  and  extracting  with  water.  It 
crystallizes  from  the  solution  in  water  in  large  rhombic 
prisms  of  the  composition  ZnSO4  -|-  7HQO.  Like  mag- 
nesium sulphate,  it  easily  loses  six  molecules  of  water, 
but  the  last  one  is  removed  with  difficulty.  It  appears, 
therefore,  that  the  constitution  of  the  salt  should  be  ex- 

roH 

OTT 

pressed  by  the  formula  SO  \  n         .    Zinc  sulphate,  as 

[o>Zn 

has  been  stated  (see  p.  596),  is  commonly  called  white 
vitriol.  It  is  easily  reduced  when  heated  with  charcoal. 
The  salt  is  used  extensively  in  the  preparation  of  cotton- 
prints  and  in  medicine. 

Zinc  Carbonate,  ZnCO3,  occurs  in  nature  as  smithson- 
ite.  The  precipitate  formed  by  adding  a  solution  of  a 
soluble  carbonate  to  a  solution  of  a  zinc  salt  is  generally 
a  basic  carbonate,  but  the  composition  varies  according  to 
the  conditions.  Dilute  solutions  of  sodium  carbonate  and 

Zn<gH 

>co 

zinc  sulphate  give  mainly  the  compound  Zn<^ 


.    Zn<OH 

or  Zn(OH)2.ZnCO3.ZnO.  Much  more  complicated  salts 
are,  however,  usually  obtained.  With  ammonia,  zinc 
carbonate  forms  a  soluble  compound. 

Reactions  which  are  of  Special  Value  in  Chemical  Anal- 
ysis. —  The  principal  reactions  which  are  made  use  of 
for  the  purpose  of  separating  zinc  from  other  elements 
have  been  mentioned  above.  These  are  the  reactions 


622  INORGANIC  CHEMISTRY. 

with  hydrogen  sulphide  and  ammonium  sulphide ;  with 
potassium  and  sodium  hydroxide ;  with  potassium  and 
sodium  carbonates  ;  and  with  ammonium  carbonate.  An- 
other reaction  which  is  used  in  analysis  is  that  which 
takes  place  when  zinc  salts  are  heated  on  charcoal  before 
the  blowpipe,  moistened  with  cobalt  nitrate,  and  ignited. 
Under  these  circumstances  a  green-colored  mass  known 
as  Rinmann's  green  is  formed,  which  is  probably  a  zin- 

cate  of  cobalt,  Zn<Q>Co. 

CADMIUM,  Cd  (At.  Wt.  111.10). 

General. — The  compounds  of  cadmium  are  very  similar 
to  those  of  zinc  and  magnesium.  The  element  occurs  in 
nature  in  much  smaller  quantity  than  either  of  these,, 
frequently  in  company  with  zinc,  and  its  compounds  ara 
not  as  frequently  met  with.  It  is  always  bivalent.  A 
mineral  known  as  greenockite  is  cadmium  sulphide,. 
CdS. 

Preparation  and  Properties. — Cadmium  is  obtained 
principally  from  different  varieties  of  zinc  blende,  and 
separates  with  the  zinc.  Being  more  volatile  than  zinc, 
it  passes  over  first  when  the  mixture  is  distilled.  From 
this  first  distillate,  which  contains  the  oxides  of  zinc  and 
cadmium,  the  metals  are  reduced  by  heating  with  char- 
coal. It  has  a  color  like  that  of  tin,  and  is  harder  than 
tin.  According  to  the  specific  gravity  of  its  vapor,  its 
molecule  is  identical  with  its  atom,  for  the  molecular 
weight  is  approximately  111. 

Cadmium  chloride,  CdCl2,  like  zinc  chloride,  is  volatile  ; 
the  sulphate  crystallizes  well,  but  is  not  analogous  in 
composition  to  the  sulphates  of  magnesium  and  zinc,  as 
the  composition  of  the  crystallized  salt  is  represented 
by  the  formula  3CdSO4  -f-  8H2O  ;  the  normal  carbonate, 
CdCO3,  is  precipitated  by  soluble  carbonates. 

Cadmium  Sulphide,  CdS,  is  one  of  the  most  character- 
istic compounds  of  the  element.  It  is  a  beautiful  yellow 
substance,  which  is  thrown  down  from  a  solution  of  a 
cadmium  salt  by  hydrogen  sulphide.  While  it  dissolves 


CADMIUM— MERCURY.  G23 

in  concentrated  acids  it  does  not  dissolve  in  dilute  acids, 
and  it  is  therefore  completely  precipitated  by  hydrogen 
sulphide.  It  is  used  as  a  constituent  of  yellow  paints. 

Cadmium  Cyanide,  Cd(CN)2,  is  formed  as  a  white  pre- 
cipitate when  potassium  cyanide  is  added  to  a  fairly 
concentrated  solution  of  a  cadmium  salt.  It  dissolves 
in  an  excess  of  potassium  cyanide  in  consequence  of  the 
formation  of  the  compound  K2Cd(CN)4. 

Analytical  Reactions.  —  Cadmium,  as  has  just  been 
stated,  is  precipitated  by  hydrogen  sulphide.  It  is 
thrown  down  together  with  the  other  elements  of  the 
hydrogen  sulphide  group  (see  p.  198).  As  the  sulphide 
is  not  soluble  in  ammonium  sulphide,  it  is  easily  sepa- 
rated from  those  of  arsenic,  antimony,  and  tin  by  treating 
with  this  reagent,  when  it  is  left  undissolved  in  company 
with  the  sulphides  of  mercury,  lead,  bismuth,  and  copper. 
The  double  salt  of  cuprous  cyanide  and  potassium  cya- 
nide is  not  decomposed  by  hydrogen  sulphide,  whereas 
the  corresponding  salt  of  cadmium  is  decomposed  by  it, 
and  the  yellow  sulphide  is  precipitated. 

The  hydroxide  of  cadmium  differs  from  that  of  zinc  in 
not  having  acid  properties.  It  does  not  dissolve  in  the 
caustic  alkalies. 

MEKCUKY,  Hg  (At.  "Wt.  198.49). 

General. — As  already  stated,  mercury  yields  two  series 
of  compounds,  known  as  mercurous  and  mercuric  com- 
pounds, which  are  analogous  to  the  two  series  of  copper 
compounds.  While,  however,  copper  forms  with  the 
oxygen  acids  only  such  salts  as  belong  to  the  cupric 
series,  as  CuSO4,  Cu(NO3)2,  etc.,  mercury  forms  salts  be- 
longing to  both  series.  There  is,  for  example,  a  mer- 
curous nitrate,  HgNO3,  and  a  mercuric  nitrate,  Hg(NO3)2 ; 
a  mercurous  sulphate,  Hg2SO4,  and  a  mercuric  sulphate, 
HgSO4,  etc.  The  mercurous  compounds  are  readily  con- 
verted into  the  mercuric  compounds  by  the  action  of 
oxidizing  agents,  and  the  mercuric  are  converted  into 
mercurous  compounds  by  the  action  of  reducing  agents. 
The  action  will  be  treated  of  under  the  individual  com- 
pounds. The  question  as  to  the  correct  formula  of  the 


624  INORGANIC  CHEMISTHY. 

rnercurous  salts  is  in  the  same  condition  as  that  in  regard 
to  the  formula  of  cuprous  salts,  with  this  difference  :  the 
molecular  weight  of  mercurous  chloride  leads  to  the 
formula  HgCl,  but  there  is  evidence  that  when  the  chlo- 
ride is  heated  some  mercury  is  set  free,  and  this  has  led 
to  the  suggestion  that  the  molecule  corresponds  to  the 
formula  Hg2Cl2,  and  that  the  compound  breaks  down 
into  mercury  and  mercuric  chloride  when  heated.  It  is, 
however,  quite  possible  that  the  compound  has  the  sim- 
pler formula,  and  that  this  under  the  influence  of  heat 
is  partly  decomposed,  as  represented  in  the  equation 

2HgCl  :=  HgCl,  +  Hg. 

The  fact  that  mercury  is  set  free  is,  therefore,  by  no 
means  satisfactory  evidence  that  the  formula  of  mer- 
curous chloride  is  Hg2Cl2,  and  in  the  present  state  of 
the  inquiry  it  is  perfectly  justifiable  to  write  the  formula 
HgCl. 

Forms  in  which  Mercury  occurs  in  Nature. — Mercury 
occurs  native  to  some  extent,  but  principally  in  the 
form  of  the  sulphide,  HgS,  which  is  known  as  cinnabar. 
This  is  sometimes  found  crystallized,  but  generally  amor- 
phous. The  chief  localities  are  Idria,  Almaden  in  Spain, 
and  New  Almaden  in  California. 

Metallurgy  of  Mercury. — In  order  to  obtain  mercury 
from  the  sulphide  this  is  roasted  in  vessels  so  constructed 
as  to  condense  and  collect  the  vapor  of  mercury  given 
off.  In  the  roasting  process  the  sulphur  is  oxidized  to 
sulphur  dioxide,  which  of  course  escapes.  In  some 
places  the  ore  is  mixed  with  limestone  and  distilled  from 
clay  or  iron  retorts,  when  the  mercury  passes  over. 
Crude  mercury  is  redistilled  in  order  to  purify  it.  It  is 
also  purified  by  treating  it  with  dilute  nitric  acid  or  with 
a  solution  of  ferric  chloride. 

Properties. — Mercury  is  a  silver-white  metal  of  a  high 
lustre.  At  ordinary  temperatures  it  is  liquid,  though  at 
—39.5°  it  becomes  solid.  Its  specific  gravity  is  13.5959. 
It  does  not  change  in  the  air  at  ordinary  temperatures. 
It  boils  at  357.25°,  and  is  converted  into  a  colorless  vapor, 
the  specific  gravity  of  which  leads  to  the  conclusion  that, 


AMALGAMS.  625 

as  in  the  case  of  cadmium,  the  molecule  and  atom  are 
identical,  or  that  the  molecule  consists  of  only  one  atom. 
It  is  insoluble  in  hydrochloric  acid  and  in  cold  sulphuric 
acid ;  but  dissolves  in  hot  concentrated  sulphuric  acid, 
and  is  easily  soluble  in  nitric  acid.  The  vapor  of  mer- 
cury is  very  poisonous. 

Applications.  —  Mercury  is  extensively  used  in  the 
manufacture  of  thermometers,  barometers,  etc.  ;  as  tin- 
amalgam  for  mirrors ;  and  in  the  processes  by  which 
gold  and  silver  are  obtained  from  their  ores. 

Amalgams. — The  alloys  which  mercury  forms  with  other 
metals  are  called  amalgams.  These  compounds  are  gen- 
erally obtained  without  difficulty  simply  by  bringing 
mercury  in  contact  with  other  metals.  Among  the 
amalgams  which  are  of  chief  interest  are  those  of  sodium, 
ammonium,  silver,  and  gold.  Sodium  amalgam  is  made  by 
bringing  mercury  and  sodium  together.  A  crystallized 
amalgam  containing  the  constituents  in  the  proportions 
represented  in  the  formula  Hg6Na  has  been  obtained. 
Generally,  sodium  amalgam  is  easily  decomposed  by 
water,  the  mercury  separating  in  the  free  state  and  the 
sodium  acting  upon  the  water,  forming  hydrogen  and 
sodium  hydroxide.  It  is  much  used  in  the  laboratory  as 
a  convenient  means  of  producing  hydrogen  in  alkaline 
solutions.  It  serves  as  an  excellent  reducing  agent  in 
some  cases.  Ammonium  amalgam  has  already  been 
spoken  of  under  the  head  of  Ammonia  (which  see).  It 
is  a  curious  substance,  which  is  formed  when  an  electric 
current  acts  upon  a  solution  of  ammonia  containing  some 
mercury  which  is  connected  with  the  negative  pole,  and 
also  very  easily  by  pouring  a  solution  of  ammonium 
chloride  upon  sodium  amalgam.  In  the  latter  case 
sodium  chloride  and  ammonium  amalgam  are  formed. 
Apparently  the  reaction  takes  place  in  accordance  with 
the  following  equation  : 

NH4C1  +  NaHg  =  NaCl  +  NH4Hg. 

The   product  is  extremely  voluminous,  and   swells   up 
during  the  reaction,  so  that  it  occupies  under  favorable 


626  INORGANIC  CHEMISTRY. 

circumstances  about  twenty  times  the  volume  occupied 
by  the  sodium  amalgam.  It  has  a  metallic  lustre,  resem- 
bling in  general  the  other  amalgams.  It  is  very  unstable 
at  the  ordinary  temperature,  breaking  down  into  mercury, 
hydrogen,  and  ammonia.  At  a  low  temperature,  how- 
ever, it  has  been  obtained  in  crystallized  form.  The 
metallic  lustre  and  general  outward  appearance  of  the 
compound  suggests  that  whatever  is  in  combination  with 
mercury  in  it  has  probably  metallic  properties,  and  this 
affords  some  confirmation  of  the  ammonium  theory,  ac- 
cording to  which  the  presence  of  the  complex,  NH4,  in 
the  salts  formed  by  ammonia  is  assumed.  Silver  amal- 
gam and  gold  amalgam  vary  in  composition  according  to 
the  method  of  preparation,  and  when  heated  are  com- 
paratively easily  decomposed. 

Mercurous  Chloride,  HgCl,  is  commonly  called  calomel. 
Like  cuprous  chloride,  CuCl,  and  argentic  chloride,  AgCl, 
it  is  insoluble  in  water.  It  is  formed  most  readily  by  re- 
ducing mercuric  chloride.  The  reduction  can  be  accom- 
plished by  means  of  sulphurous  acid,  when  the  following 
reaction  takes  place  : 

2HgCl2  +  2H20  +  S02  -  2HgCl  +  H2SO4  +  2HC1. 

It  is  also  formed  by  heating  together  mercuric  chloride 
and  mercury,  and  by  subliming  a  mixture  of  mercuric 
sulphate,  sodium  chloride,  and  mercury.  This  method 
is  the  one  mostly  used  in  the  manufacture  of  calomel. 
The  product  obtained  by  sublimation  is  crystalline  ;  the 
precipitated  substance  forms  a  loose  powder.  As  was 
stated  above,  the  specific  gravity  of  the  vapor  corre- 
sponds to  that  required  for  the  formula  HgCl.  When 
acted  upon  for  some  time  by  light  it  undergoes  partial 
decomposition  into  mercury  and  mercuric  chloride.  This 
is  a  fact  of  great  importance,  inasmuch  as  calomel  is 
much  used  in  medicine,  and  mercuric  chloride  is  an  active 
poison.  Bottles  in  which  calomel  is  kept  should  be  care- 
fully protected  from  the  action  of  the  light. 

Just  as  mercuric  chloride  is  converted  into  mercurous 
chloride  by  reducing  agents,  so  the  latter  is  converted 
into  the  former  by  oxidizing  agents.  When,  for  example, 


MERCURIC  CHLORIDE.  627 

mercurous  chloride  is  treated  with  nitric  acid  it  is  con- 
verted into  mercuric  chloride  and  mercuric  nitrate,  as 
represented  in  the  equation 

6HgCl  +  8HNO3  =  3Hg(NO3)2  +  3HgCl3  +  2NO  +  4H2O. 

If  hydrochloric  acid  is  present  in  sufficient  quantity  the 
action  takes  place  thus  : 

3HgCl  +  3HC1  +  HNO3  =  3HgCl2  +  2H3O  +  NO. 

Further,  the  conversion  of  mercurous  nitrate  into  mer- 
curic nitrate  is  represented  by  the  equation 

3HgNO3  +  4HN03  =  3Hg(NO3)2  +  2H2O  +  NO. 

Finally,  the  action  of  oxidizing  agents  in  general  upon 
mercurous  chloride  in  the  presence  of  hydrochloric  acid 
takes  place  thus : 

2HgCl  +  2HC1  +  O  =  2HgCl2  +  H2O. 

Similar  transformations  take  place  by  treating  ferrous, 
stannous,  and  manganous  compounds  with  oxidizing 
agents  ;  and  they  will  be  taken  up  farther  on. 

Mercuric  Chloride,  or  Corrosive  Sublimate,  HgCl2,  which 
is  made  by  subliming  a  mixture  of  sodium  chloride  and 
mercuric  sulphate, 

HgS04  +  2NaCl  =  HgCl2  +  Na.SO4, 

and  by  dissolving  mercury  in  aqua  regia,  evaporating  to 
dryness,  and  subliming  the  residue,  is  a  white,  trans- 
parent, crystalline  mass,  which  is  soluble  in  water,  and 
can  be  obtained  in  crystalline  form  from  the  solution. 
It  is  more  easily  soluble  in  alcohol  and  ether  than  in 
water,  and  is  extracted  from  a  water  solution  by  shaking 
with  ether.  It  is  quite  volatile,  and  the  specific  gravity 
of  its  vapor  corresponds  to  that  required  for  the  formula 
HgCl2.  It  is  easily  reduced  to  mercurous  chloride  by 


628  INORGANIC  CHEMISTRY. 

contact  with  organic  substances,  and  by  reducing  agents 
in  general.  The  action  of  sulphur  dioxide  has  already 
been  treated  of  as  furnishing  a  method  for  the  prepara- 
tion of  mercurous  chloride.  Stannous  chloride  abstracts 
chlorine  from  it  and  forms  mercurous  chloride  and  me- 
tallic mercury,  while  the  stannous  chloride  is  converted 
into  stannic  chloride  : 

2HgCl2  +  SnCl2  =  2HgCl  +  SnCl4  ; 
HgCl2  +  SnCl2  =  Hg       +  SnCl4. 

Mercuric  chloride  is  an  active  poison,  and  has  been 
used  extensively  in  this  capacity.  It  has  a  very  marked 
influence  upon  the  lower  organisms,  which  play  such  an 
important  part  in  producing  disease  and  the  decay  of 
organic  substances,  and  is  used  as  a  disinfectant.  Wood 
impregnated  with  a  solution  of  it  is  partly  protected  from 
decay.  In  surgery  it  is  used  for  the  purpose  of  pre- 
venting contamination  of  wounds  by  the  hands  and  in- 
struments of  the  surgeon,  it  being  customary  now  for  the 
surgeon  to  wash  his  hands  and  instruments  in  a  dilute 
solution  of  the  chloride  before  performing  an  operation. 

Mercuric  chloride  unites  with  other  chlorides,  forming 
well-characterized  double  chlorides,  or  chlor-mercurates, 
which  are  analogous  to  the  double  chlorides  of  magne- 
sium, zinc,  etc.  Three  potassium  salts  are  known,  KHgCl3, 

K2HgCl4,    and    KHg2Cl5;   or  Hg<g^  Hg<g|,  and 

H   <C1 

C12       .     Similar  salts  are  formed  with  the  chlorides 


of  sodium,  ammonium,  calcium,  barium,  strontium,  and 
other  metals.  Further,  hydrochloric  acid  forms,  with 
mercuric  chloride,  a  crystallized  compound  of  the  for- 
mula HHg2Cl6  or  2HgCl2.HCl,  which  is  plainly  the  acid 
from  which  the  potassium  salt  KHg2ClB  is  derived. 

Mercurous  Iodide,  Hgl,  can  be  made  by  treating  mer- 
cury with  iodine  ;  or  by  treating  mercuric  iodide  with 
mercury  ;  and  more  easily  by  adding  potassium  iodide 
to  a  solution  of  a  mercurous  salt,  when  it  is  thrown  down 


MERCURIC  IODIDE.  629 

as' a  greenish-yellow  powder.  It  is  unstable,  and  breaks 
down,  yielding  mercury  and  mercuric  iodide.  When 
treated  with  potassium  iodide  it  suffers  the  same  de- 
composition, the  iodide  which  is  formed  combining  with 
the  potassium  iodide,  while  mercury  is  deposited.  Mer- 
curous  iodide  is  used  in  medicine.  As  it  is  more  easily 
decomposed  than  mercurous  chloride,  and  mercuric 
iodide  is  poisonous,  care  should  be  taken  in  its  use. 

Mercuric  Iodide,  HgI2,  is  made  by  direct  combination 
of  mercury  and  iodine,  and  by  the  action  of  potassium 
iodide  upon  a  solution  of  a  mercuric  salt,  when  it  is  pre- 
cipitated as  a  beautiful  scarlet-red  powder.  Though 
insoluble  in  water,  it  is  soluble  in  alcohol  and  ether.  It 
dissolves,  further,  in  a  solution  of  potassium  iodide  in 
consequence  of  the  formation  of  a  double  iodide,  or  iodo- 
mercurate.  When  heated  to  about  150°  the  color 
changes  from  red  to  yellow,  and  after  cooling  the  yel- 
low substance  changes  to  red.  Sometimes  it  can  be  kept 
in  the  condition  in  which  it  has  a  yellow  color  for  some 
time  after  it  has  cooled  down,  and  then  if  touched  at  one 
point  the  change  to  the  red  substance  takes  place  rap- 
idly through  the  entire  mass.  Both  the  red  and  the  yel- 
low modifications  are  crystallized,  but  in  different  forms. 
The  red  crystals  are  tetragonal ;  the  yellow  ones  are 
rhombic  or  monoclinic.  The  monoclinic  structure  is 
evidently,  from  what  has  been  said,  an  unstable  one  for 
this  compound.  It  seems  probable  that  the  difference 
between  the  two  forms  is  due  to  a  difference  in  the  com- 
plexity of  the  molecules  ;  perhaps  the  larger  molecules  of 
the  red  iodide  are  decomposed  by  heat  into  the  smaller 
molecules  of  the  yellow  iodide.  When  the  iodide  is  first 
precipitated  from  a  solution  of  mercuric  chloride  by 
potassium  iodide  it  is  yellow,  but  it  rapidly  turns  red. 

Just  as  mercuric  chloride  forms  a  compound  with  hy- 
drochloric acid,  so  mercuric  iodide  forms  similar  com- 
pounds. These  have  the  composition  represented  by 
the  formulas  HgI2.HI  or  HHgI3,  and  2HgI2.3HI  or 
H3Hg2I7.  It  also  combines  with  iodine,  forming  the  per- 
iodide  HgI9,  which  is  decomposed  by  water,  forming  mer- 
curic iodide.  With  potassium  chloride  it  forms  the  salts 


G30  INORGANIC  CHEMISTRY. 

K2HgI4  or  HgI2.2KI,  and  KHgI3  or  HgIa.KI.  These  salts 
form  colorless  solutions.  They  are  also  formed  by  the 
action  of  potassium  iodide  upon  mercurous  iodide,  when 
mercury  separates : 

2HgI  +  2KI  =  K^Hgl,  +  Hg. 

Similar  double  iodides  are  formed  with  the  chlorides  of 
sodium,  ammonium,  barium,  calcium,  strontium,  magne- 
sium, and  other  metals. 

Mercurous  Oxide,  Hg2O,  is  formed  by  treating  a  mer- 
curous salt  with  potassium  hydroxide.  It  is  a  black 
powder,  and  decomposes  easily  into  mercuric  oxide  and 
mercury  when  the  light  shines  upon  it. 

Mercuric  Oxide,  HgO,  like  the  iodide,  presents  itself  in 
two  colors — the  red  and  the  yellow.  When  mercury  is 
heated  for  a  long  time  at  a  temperature  near  the  boiling- 
point  so  that  the  air  has  access  to  it,  it  is  converted  into 
the  red  oxide.  When  a  solution  of  a  mercuric  salt  is 
treated  with  caustic  soda  or  caustic  potash  the  yellow 
oxide  is  precipitated.  When  heated,  the  red  oxide  be- 
comes darker,  and  finally  nearly  black,  and  on  cooling 
the  red  color  reappears.  The  yellow  oxide  also  becomes 
darker  on  heating.  Exposed  to  the  light,  the  red  oxide 
loses  some  of  its  oxygen,  and  mercury  is  deposited. 
When  heated  to  a  sufficiently  high  temperature  both 
varieties  give  up  all  their  oxygen,  as  was  seen  in  the 
case  of  the  red  oxide  in  the  preparation  of  oxygen.  It 
will  be  remembered  that  oxygen  was  discovered  by  heat- 
ing this  compound.  Of  what  overwhelming  importance 
this  discovery  was  the  student  will  now  better  appreci- 
ate than  when  the  discovery  was  first  mentioned.  It  will 
now  be  seen  that  most  of  the  chemical  phenomena  with 
which  we  have  to  deal  involve  the  action  of  oxygen. 

Hydroxides  of  mercury  are  not  known. 

Mercurous  Sulphide,  Hg2S. — There  seems  to  be  some 
question  whether  a  compound  of  the  formula  Hg2S 
exists  or  not.  According  to  the  latest  investigations 
on  the  subject,  the  precipitate  which  is  formed  when 
hydrogen  sulphide  is  passed  into  a  solution  of  a  nier- 


MERCURIC  SULPHIDE.  631. 

curous  salt  is  a  mixture  of  mercuric  sulphide,  HgS, 
and  mercury.  At  all  events,  it  is  certain  that,  even 
if  mercurous  sulphide  exists,  it  breaks  down  with  great 
ease  into  the  mercuric  compound  and  mercury. 

Mercuric  Sulphide,  HgS,  is  the  principal  compound  of 
mercury  found  in  nature.  This  natural  variety,  which 
is  known  as  cinnabar,  is  a  red  crystallized  compound. 
When  prepared  by  treating  a  solution  of  a  mercuric  salt 
with  hydrogen  sulphide  it  is  an  amorphous  black  pow- 
der. This  black  powder  can,  however,  be  converted  into 
the  red  crystallized  variety  by  sublimation,  and  by  allow- 
ing it  to  stand  in  contact  with  a  solution  of  the  sulphide 
of  an  alkali  metal.  The  sulphide,  whether  amorphous 
or  crystallized,  is  acted  upon  with  difficulty  by  acids. 
Dilute  nitric  acid,  which  easily  dissolves  the  sulphides 
of  the  other  metals  which  resemble  mercury,  leaves  it 
unacted  upon,  and  advantage  is  taken  of  this  -fact  for  the 
purpose  of  separating  it  from  the  sulphides  of  lead,  bis- 
muth, copper,  and  cadmium  in  qualitative  analysis.  It  is 
dissolved  by  concentrated  nitric  acid  and  by  aqua  regia. 
In  the  former  case  a  white  insoluble  compound  of  mer- 
curic nitrate  and  mercuric  sulphide,  Hg(NO3)2.2HgS, 
is  sometimes  formed.  Potassium  hydrosulphide  dis- 
solves mercuric  sulphide,  forming  a  compound,  K2HgS2, 
or  K2S.HgS,  which  is  easily  decomposed  by  water,  mer- 
curic sulphide  being  thrown  down.  The  substance  used 
in  medicine  under  the  name  EtMops  martialis  is  an  inti- 
mate mixture  of  amorphous  sulphide  and  sulphur.  Like 
the  oxide,  cinnabar  turns  dark  when  heated,  and  it  slowly 
undergoes  the  same  change  when  exposed  to  the  light,  in 
consequence  of  a  slight  decomposition  into  mercury  and 
sulphur.  When  that  which  has  not  been  heated  to  too 
high  a  temperature  is  cooled  down  again,  it  acquires  its 
original  red  color.  If  it  has  been  heated  to  the  tempera- 
ture of  sublimation,  the  black  color  of  that  which  does 
not  sublime  is  permanent.  Cinnabar  is  used  as  a  pig- 
ment in  the  manufacture  of  red  paints. 

Mercuric  Cyanide,  Hg(CN)2,  is  formed  by  dissolving 
mercuric  oxide  in  an  aqueous  solution  of  hydrocyanic 


632  INORGANIC  CHEMISTRY. 

acid.  It  is  soluble  in  water,  and  crystallizes  well  in 
quadratic  prisms.  When  heated  it  is  decomposed  into 
cyanogen  and  mercury,  and  this  affords  the  most  con- 
venient method  of  making  cyanogen.  When  heated  rap- 
idly there  is  formed  a  considerable  quantity  of  a  black 
substance  called  paracyanogen,  which  has  the  same  per- 
centage composition  as  cyanogen,  but  is  not  volatile.  This 
is  probably  a  polymeric  variety  of  cyanogen.  The 
cyanide  unites  with  the  cyanides  of  other  metals,  form- 
ing double  cyanides,  of  which  those  corresponding  to 
the  following  formulas  are  good  examples  :  K2Hg(CN)4, 
MgHg(CN)4,  etc.  These  are  soluble  in  water,  and,  as  a 
rule,  crystallize  well. 

Fulminating  mercury  is  an  explosive  compound  much 
used  in  the  manufacture  of  gun-caps.  It  is  made  by  dis- 
solving mercury  in  nitric  acid  and  adding  alcohol.  Its 
explosion  consists  in  a  sudden  breaking  down  into  nitro- 
gen, carbon  dioxide,  and  mercury.  The  composition  of 
the  compound  is  represented  by  the  formula  C2N2O2Hg. 
It  would  lead  too  far  to  discuss  its  constitution  —  a  sub- 
ject which  has  occupied  the  attention  of  some  of  the 
most  celebrated  chemists. 

Mercurous  Nitrate,  HgNO3,  is  made  by  treating  nitric 
acid  with  an  excess  of  mercury.  This  salt  is  easily  de- 
composed, forming  difficultly  soluble  basic  salts,  some  of 
which  are  of  complicated  composition.  The  simplest  is 
that  which  has  the  composition  HgOH.HgNO3.  This 
is  easily  explained  on  the  assumption  that  the  nitrate 

Hg(NO3) 
has  the  formula  I  The  decomposition  by  water 

is  then  represented  by  the  equation 


a+H20  = 
Hg(N03)  Hg(OH) 


This  may  possibly  be  regarded  as  a  slight  argument  in 
favor  of  the  doubled  formula  for  the  mercurous  com- 
pound. It  is,  however,  far  from  conclusive,  as  the  salt 
may  be  considered  to  be  made  up  as  represented  in  this 


COMPOUNDS  OF  SALTS  OF  MERCURY  AND  AMMONIA.  633 

(OH 

formula,  ON  •<  OHg ,  according  to  which  it  is  a  derivative 
(OHg 

of  the  acid  NO(OH)3,  corresponding  to  phosphoric  acid, 
PO(OH)3. 

Mercuric  Nitrate,  Hg(NO3)2,  is  formed  by  treating  mer- 
curic oxide  with  an  excess  of  nitric  acid  and  evaporating, 
when  the  salt  can,  under  favorable  circumstances,  be 
obtained  in  crystals.  It  is  easily  decomposed  by  water, 
with  formation  of  a  basic  salt  which  is  insoluble  in 
water.  It  has  the  composition  represented  by  the  for- 
mula NO3-Hg-O-Hg-O-Hg-NO3  +  H2O. 

Compounds  formed  by  Salts  of  Mercury  with  Ammonia. 
— Attention  has  already  been  called  to  the  compounds 
formed  by  the  salts  of  copper  with  ammonia,  and  the 
statement  was  made  that  this  power  to  combine  with  am- 
monia and  with  ammonium  salts  is  very  common  among 
the  salts.  Some  of  these  compounds  which  are  formed 
by  the  salts  of  mercury  are  of  special  interest.  When 
mercuric  chloride  is  treated  with  ammonia  a  white  pre- 
cipitate is  formed,  the  composition  of  which  is  repre- 
sented by  the  formula  HgClNH2.  The  formation  takes 
place  according  to  the  equation 

HgCl2  +  2NH3  =  HgClNH2  +  NH4C1. 

The  simplest  view  in  regard  to  the  constitution  of  this 
compound  is  that  it  is  mercuric  chloride  in  which  a 
chlorine  atom  is  replaced  by  the  group  NH2,  or  amide, 

Cl 
as  represented  in  the  formula  Hg<^-rr  .     According  to 

this  view  the  compound  is  called  mercuric  chloramide. 
It  is  known  as  white  precipitate,  or,  to  distinguish  it  from 
another  similar  compound,  as  infusible  white  precipitate. 
It  is  also  possible  that  the  constitution  of  the  compound 

(H2 
should  be  represented  by  the  formula  N<  Hg,  according 

(a 

to  which  it  is  ammonium  chloride  in  which  two  atoms  of 
hydrogen  are  replaced  by  an  atom  of  bivalent  mercury. 
The  similar  compound  referred  to  is  formed  by  adding  a 
solution  of  mercuric  chloride  to  a  boiling  solution  of  am- 


634  INORGANIC  CHEMISTRY. 

monium  chloride  containing  ammonia,  and  separates  out 
on  cooling.  It  has  the  composition  Hg(NH3Cl)2,  and  is 
to  be  regarded  as  ammonium  chloride,  in  two  molecules 
of  which  two  hydrogen  atoms  are  replaced  by  an  atom 
of  bivalent  mercury,  as  represented  in  the  formula 

This  is  called  mercuric  diammonium  chlo- 


ride, or  fusible  white  precipitate.  Mercurous  compounds 
corresponding  to  both  the  above-mentioned  derivatives 
of  mercuric  chloride  are  known.  The  first,  or  mercurous 
chloramide,  Hg2NH2Cl,  is  the  black  substance  formed 
when  mercurous  chloride  is  treated  with  ammonia  : 

2HgCl  +  2NH3  =  Hg2NH2Cl  +  NH4C1. 

The  second,  or  mercurous  ammonium  chloride,  HgNH3Cl, 
is  formed  by  treating  calomel  with  ammonia  gas. 

Similar  compounds  are  obtained  from  the  other  salts  of 
mercury.  In  a  solution  of  mercurous  nitrate  ammonia 
forms  a  black  precipitate,  known  in  pharmacy  as  Mercu- 
rius  solubilis  Hahnemanni.  It  may  be  regarded  as  derived 
from  ammonium  nitrate  by  replacement  of  two  hydrogen 
atoms  by  two  atoms  of  mercury,  as  represented  in  the 
formula  Hg2H2N.NO3. 

Reactions  which  are  of  Special  Value  in  Chemical 
Analysis.  —  As  has  been  seen  in  the  account  already  given 
of  the  conduct  of  the  compounds  of  mercury,  mercurous 
and  mercuric  compounds  conduct  themselves  quite  dif- 
ferently. Among  the  most  characteristic  reactions  of 
mercurous  compounds  are  the  following  : 

Sodium  or  potassium  hydroxide  gives  a  black  precipi- 
tate of  mercurous  oxide. 

Ammonia  gives  a  black  precipitate  which  is  a  compound 
of  mercurous  oxide  and  ammonia. 

Hydrochloric  acid  and  soluble  chlorides  form  mercurous. 
chloride,  and  the  precipitate  turns  black  when  treated 
with  ammonia. 

The  reactions  with  stannous  chloride,  potassium  iodide, 
hydrogen  sulphide,  and  ammonium  sulphide  have  been 
explained. 

The  principal  reactions  of  the  mercuric  salts  are  the 


GALLIUM.  635 

following :    Sodium  or  potassium  hydroxide  produces  a 
yellow  precipitate  of  mercuric  oxide. 

The  reactions  with  hydrogen  sulphide,  stannous  chlo- 
ride, and  potassium  iodide  have  been  described  above. 

ELEMENTS  OF  FAMILY  III,  GROUP  B: 
GALLIUM— INDIUM— THALLIUM. 

General. — All  the  elements  of  this  group  are  rare. 
Gallium  forms  compounds  which  in  composition  are 
analogous  to  those  of  aluminium,  in  which  it  is  trivalent, 
as  in  the  compounds  Ga013,  Ga2O3,  Ga(NO3)3,  Ga2(SO4)3, 
etc.  On  the  other  hand,  it  also  forms  compounds  in  which 
it  is  bivalent,  as  in  the  chloride  GaCl2  and  GaO.  Indium 
also  forms  compounds  in  which  it  is  trivalent,  and  a  few 
in  which  it  is  bivalent.  Thallium,  like  gold,  is  trivalent 
and  univalent  in  its  compounds. 

Gallium,  Ga  (At.  Wt.  69.38). — This  element  is  found  in 
some  varieties  of  zinc  blende,  and  was  discovered  in  that 
which  occurs  at  Pierrefitte,  in  France.  It  owes  its  name 
to  the  Latin  name  of  France,  Gallia.  Like  scandium,  it 
is  of  special  interest  for  the  reason  that  its  properties 
were  foretold  by  Mendeleeff  some  years  before  it  was  dis- 
covered, and  it  was  described  by  him  under  the  name  of 
eka-aluminium,  as  scandium  was  described  under  the 
name  of  eka-boron.  The  metal  has  a  bluish- white  color, 
and  does  not  lose  its  lustre  in  the  air.  It  does  not  de- 
compose water.  Even  when  heated  to  a  high  tempera- 
ture in  oxygen  it  only  becomes  covered  with  a  thin  layer 
of  oxide.  It  dissolves  in  hydrochloric  acid  and  in  po- 
tassium hydroxide  with  evolution  of  hydrogen. 

Compounds  of  Gallium. — The  chlorides  of  gallium, 
GaCl2  and  GaCl8,  are  formed  by  treating  the  metal  with 
chlorine.  With  water  both  give  basic  chlorides.  That 
formed  from  gallons  chloride,  GaCl2,  is  completely  con- 
verted into  gallium  oxide  by  further  action  of  water. 
Gallic  sulphate,  Ga2(SO4)3,  is  decomposed  by  boiling  water, 
a  basic  salt  being  formed.  With  ammonium  sulphate  it 
forms  a  double  sulphate  of  the  formula  NH4Ga(SO4)2  -f- 
12H2O,  which  is  perfectly  analogous  to  ammonium  alum 
(see  Alums,  p.  576). 


636  INORGANIC  CHEMISTRY. 

Indium,  In  (At.  Wt.  112.99). — Indium  was  discovered  in 
a  variety  of  zinc  blende  found  at  Freiberg,  in  Germany. 
The  discovery  was  made  by  means  of  the  spectroscope, 
and  as  the  spectrum  of  the  element  contains  two  very 
characteristic  blue  lines  it  was  called  indium.  It  has 
since  been  found  in  other  varieties  of  zinc  blende,  but 
always  in  small  quantity.  It  is  a  soft,  white,  lustrous 
metal.  It  does  not  undergo  change  in  contact  with  the 
air  at  ordinary  temperatures,  but  when  heated  it  takes 
fire  and  burns,  forming  the  oxide.  It  does  not  decom- 
pose water  even  at  the  boiling  temperature. 

Compounds  of  Indium. — With  chlorine  indium  forms 
the  compound  InCl3,  which  is  volatile  at  a  high  tempera- 
ture. The  specific  gravity  of  the  vapor  is  that  required 
for  the  formula  InCl3.  The  sulphate  is  easily  formed 
by  dissolving  the  metal  in  sulphuric  acid.  It  combines 
with  ammonium  sulphate  to  form  the  double  sulphate 
]STH4In(SO4)2  -f  12H2O,  analogous  to  alum.  The  corre- 
sponding potassium  salt  does  not  contain  the  same  num- 
ber of  molecules  of  water  of  crystallization. 

Thallium,  Tl  (At,  Wt.  202.61).— This  element  was  dis- 
covered in  the  flue-dust  of  a  sulphuric-acid  factory  in  the 
Harz  mountains,  by  the  aid  of  the  spectroscope.  As  it 
colors  the  flame  a  beautiful  green  it  was  called  thallium, 
from  the  Greek  $orAAo£,  which  signifies  a  green  branch. 
It  has  since  been  found  in  a  number  of  varieties  of  iron 
pyrites  and  copper  pyrites.  It  is  a  soft,  bluish-white 
metal,  and  has  a  lustre  like  lead.  It  is  oxidized  when 
heated  to  a  sufficiently  high  temperature  in  the  air. 

Compounds  of  Thallium. — When  thallium  is  treated 
with  chlorine  it  is  converted  into  thallous  chloride,  T1C1 ; 
and  when  this  is  treated  under  water  with  chlorine  it  is 
converted  into  the  trichloride  or  thallic  chloride,  T1C1,. 
Thallous  chloride,  further,  is  formed  as  a  caseous  pre- 
cipitate when  hydrochloric  acid  is  added  to  a  solution  of 
a  thallous  salt.  When  exposed  to  the  light  it  turns  violet. 
— Thallous  hydroxide,  Tl(OH),  is  formed  by  the  action  of 
the  metal  on  water  in  the  air,  and  by  treating  a  solution 
of  the  sulphate  with  barium  hydroxide.  It  is  easily 
soluble  in  water,  and  the  solution  has  an  alkaline  reac- 


COMPOUNDS  OF  THALLIUM.  037 

tion. —  Thallic  hydroxide,  T1(OH)3,  is  formed  by  treating 
a  solution  of  thallic  chloride  with  potassium  hydroxide. 
When  dried  it  loses  water  and  forms  the  compound 
T1O.OH,  analogous  to  metaboric  acid,  BO. OH,  and  meta- 
aluminic  acid,  A1O.OH.  Thallium  forms  an  insoluble 
chloride,  T1C1,  which  turns  violet  when  exposed  to  the 
light,  and  in  this  respect  it  resembles  silver  ;  the  forma- 
tion of  the  soluble  hydroxide,  Tl(OH),  which  has  an 
alkaline  reaction,  is  highly  suggestive  of  the  alkali 
metals ;  while  the  formation  of  the  hydroxide,  TIO(OH), 
shows  that  thallium  is  allied  to  aluminium.  An  exami- 
nation of  its  salts  shows  that  those  in  which  it  is  univa- 
lent  resemble  the  salts  of  the  alkali  metals,  while  those 
in  which  it  is  trivalent  resemble  the  salts  of  aluminium. 
-Thallous  sulphate,  T12SO4,  and  the  thallous  phosphates, 
H2T1PO4,  HT12PO4,  and  T13PO4,  are  isomorphous  with 
the  corresponding  salts  of  potassium. — Thallous  sulphide 
is  thrown  down  as  a  black  powder  when  hydrogen  sul- 
phide is  passed  into  a  solution  of  a  thallous  salt. — Thai' 
lie  sulphate,  T12(SO4)3,  combines  with  sulphates  of  the 
alkali  metals,  forming  salts  of  the  general  formula 
T1M(SO4)2 ;  but  these  do  not  crystallize  like  the  alums. 
On  the  other  hand,  thallous  sulphate,  T12SO4,  combines 
with  aluminium  sulphate,  and  other  similar  sulphates, 
forming  salts  perfectly  analogous  to  the  alums,  and  in 
these  the  thallium  takes  the  place  of  the  alkali  metal,  as, 
for  example,  in  the  salts 

TlAl(SO4)f  +  12HaO,     TlFe(SO4),  +  12H2O,     etc. 


CHAPTER  XXX. 

ELEMENTS  OF  FAMILY  IV,  GROUP  B  : 
GERMANIUM— TIN— LEAD. 

General. — Of  the  elements  of  this  group  germanium  is 
extremely  rare.  It  was  discovered  by  Winkler  in  1885. 
Tin  and  lead,  on  the  other  hand,  have  long  been  known, 
and  are  extensively  used.  All  form  two  series  of  com- 
pounds, in  one  of  which  they  are  bivalent  and  in  the 
other  quadrivalent.  The  general  formulas  of  some  of 
the  principal  compounds  of  the  first  series  are  as  fol- 
lows: 

MC12,  MO,  M(OH)2,  M(NO3)2,  MSO4,  MCO3,  etc. 

The  general  formulas  of  some  of  the  principal  compounds 
of  the  other  series  are  as  follows : 

MC14,  M02,  M(OH)4,  MO(OH)2,  etc. 

These  elements  have  already  been  referred  to  at  the 
close  of  Chapter  XXII.  (see  p.  425),  and  attention  was 
then  called  to  the  resemblance  between  them  and  carbon 
and  silicon.  In  this  connection  it  will  be  well  to  repeat 
what  was  there  said.  Of  the  three  elements  of  the 
group,  germanium  and  tin  are  more  acidic  in  character 
than  lead.  They  combine  with  chlorine  in  two  propor- 
tions, forming  the  chlorides  GeCl3,  SnCla,  PbCla,  GeCl4, 
SnCl4,  PbCl4.  With  oxygen  they  combine  to  form  the 
compounds  GeOa,  SnO2,  and  PbO2.  Stannic  oxide,  SnO2, 
and  lead  peroxide,  PbO2,  form  salts  with  bases,  and  these 
have  the  composition  represented  by  the  general  formulas 
M2SnO3  and  M2PbO3,  and  are  therefore  analogous  to  the 
silicates,  carbonates,  and  titanates.  On  the  other  hand, 
further,  salts  are  known  which  are  derived  from  the 
oxide  PbO.  These  have  the  general  formula  M2PbOa>. 

(638) 


GERMANIUM—  TIN.  639 

and  are  to  be  regarded  as  salts  of  an  acid  Pb(OH)2. 
These  salts  are  not  stable,  and  are  not  easily  obtained. 
Most  of  the  derivatives  of  lead  are  those  in  which  it 
plays  the  part  of  a  base-forming  element.  Notwith- 
standing the  marked  analogy  between  some  of  the  com- 
pounds of  tin  and  those  of  the  members  of  the  silicon 
group,  it  appears,  on  the  whole,  advisable  to  treat  of  this 
element  in  company  with  lead,  which  it  also  resembles  in 
many  respects. 

GERMANIUM,  Ge  (At.  Wt.  71.93). 

Germanium  is  the  third  element  the  properties  of 
which  were  foretold  by  Mendeleeff  by  the  aid  of  the 
periodic  law.  As  it  occurs  in  the  silicon  group  he  called 
it  eka-silicon.  It  was  discovered  in  a  silver  ore  occur- 
ring at  Freiberg,  Germany.  The  name  has,  of  course, 
reference  to  the  country  in  which  it  was  discovered.  It 
acts  mostly  as  a  base-forming  element,  being  perhaps 
more  like  tin  than  any  other  one  metal.  It  forms  the 
two  chlorides  GeCl2  and  GeCl4,  and  the  corresponding 
fluorides  GeF2  and  GeF4 ;  but  preferably  it  forms  those 
compounds  in  which  the  element  is  quadrivalent.  The 
fluoride  forms  double  salts  resembling  the  fluosilicates, 
as,  for  example,  pot assium fluogermanat e,  K2GeF6. 

TIN,  Sn(At.  Wt.  118.15). 

General. — The  compounds  of  tin  with  which  we  gen- 
erally have  to  deal  belong,  with  the  exception  of  stan- 
nous  chloride,  to  the  series  in  which  the  metal  is  quad- 
rivalent, and  in  this  series  it  acts  as  an  acid-forming 
element.  The  chloride,  SnCl4,  corresponds  to  the  chlo- 
rides of  carbon  and  silicon,  CC14  and  SiCl4.  Unlike 
these  elements,  however,  it  does  not  form  a  compound 
with  hydrogen. 

Occurrence. — Tin  occurs  almost  exclusively  as  tin 
stone  or  cassiterite  in  nature.  This  is  the  dioxide,  SnO2, 
corresponding  to  carbon  dioxide,  CO2 ;  silicon  dioxide, 
SiO2 ;  titanium  dioxide,  TiO2 ;  etc.  It  also  occurs  in 
small  quantities  in  company  with  gold  as  metallic  tin, 


640  INORGANIC  CHEMISTRY. 

and  in  a  variety  of  pyrites  of  the  formula  Cu4SnS4  -)- 
FeSnS4,  known  as  stannite. 

Metallurgy. — The  ores  are  roasted  for  the  purpose  of 
getting  rid  of  the  sulphur  and  arsenic,  and  the  oxide  is 
then  heated  with  coal  in  a  furnace.  After  the  reduction 
is  complete  the  tin  is  drawn  off  and  cast  in  bars.  This 
tin  is  impure,  and  when  again  slowly  melted,  that  which 
first  melts  is  purer.  By  letting  it  run  off  as  soon  as  it 
melts  the  comparatively  difficultly  fusible  alloy  remains 
behind,  and  the  tin  is  thus  rendered  much  purer.  The 
commercial  variety  of  tin  known  as  Bonca  tin  is  the 
purest.  It  receives  its  name  from  Banca,  in  the  East 
Indies,  where  it  is  made.  Block-tin  is  made  in  England, 
and  is  also  comparatively  pure. 

Properties. — Tin  is  a  white  metal,  which  in  general  ap- 
pearance resembles  silver.  It  is  soft  and  malleable,  and 
can  be  hammered  out  into  very  thin  sheets,  forming  the 
well-known  tin-foil.  Its  specific  gravity  is  7.3.  At  200° 
it  is  brittle,  and  at  228°  it  melts.  At  ordinary  tempera- 
tures it  remains  unchanged  in  the  air.  It  dissolves  in 
hydrochloric  acid,  forming  stannous  chloride,  SnCl2 ;  in 
sulphuric  acid,  forming  stannous  sulphate,  SnSO4,  sul- 
phur dioxide  being  evolved  at  the  same  time.  Ordinary 
concentrated  nitric  acid  oxidizes  it,  the  product  being  a 
compound  of  tin,  oxygen,  and  hydrogen,  known  as  meta- 
stannic  acid,  which  is  a  white,  powder  insoluble  in  ni- 
tric acid  and  in  water.  It  is  dissolved  by  a  hot  solution 
of  potassium  hydroxide  which  forms  potassium  stannate, 
K2Sn03. 

Applications. — It  is  used  in  making  alloys,  of  which 
bronze  (see  p.  591),  soft  solder,  and  britannia  metal  are 
the  most  important.  It  is  used  also  for  protecting  other 
metals,  as  in  the  tinware  vessels  in  such  common  use, 
which  are  made  of  iron  covered  with  a  layer  of  tin. 
Copper  vessels  are  also  frequently  covered  with  tin. 

Alloys. — Bronze  has  already  been  treated  of  under 
Copper.  It  is  made  of  copper,  tin,  and  zinc.  Soft  solder 
is  made  of  equal  parts  of  tin  and  lead,  or  of  two  parts  of 
tin  and  one  of  lead.  Britannia  metal  is  composed  of  nine 
parts  of  tin  and  one  of  antimony.  Tin  amalgam  is  made 


STANNOUS  CHLORIDE.  641 

by  bringing  tin  and  mercury  together,  and  is  used  in  the 
silvering  of  mirrors. 

Stannous  Chloride,  SnCl2,  is  formed  by  dissolving  tin 
in  hydrochloric  acid,  and  if  the  solution  is  concentrated 
enough  the  compound  crystallizes  oat.  The  crystals  have 
the  composition  SnCl2  -j-  2H2O.  This  is  the  commercial 
product  known  as  tin  salt.  It  is  very  easily  soluble  in 
water,  but  if  the  solution  is  dilute  it  becomes  turbid  in 
consequence  of  the  formation  of  the  insoluble  basic  salt, 


2SnCla  +  3H2O  =  2Sn<j    +  H2O  +  2HCL 

This  same  decomposition  takes  place  if  the  solution  is 
allowed  to  stand  in  contact  with  the  air.  Under  these 
circumstances  a  part  of  the  stannous  chloride  is  converted 
into  stannic  chloride  by  oxidation  : 

3SnCl2  +  H2O  +  O  =  SnCl4  +  2Sn(OH)Cl. 

"When  the  crystals  are  heated  they  melt  at  about  40°,  and 
if  heated  higher  they  undergo  partial  decomposition, 
forming  the  oxide  and  hydrochloric  acid  : 

SnCl2  +  H2O  =  SnO  +  2HC1. 

Stannous  chloride  has  a  marked  tendency  to  combine 
with  chlorine  and  to  pass  over  into  stannic  chloride. 
This  power  has  already  been  shown  in  its  action  upon 
mercuric  chloride  and  upon  mercurous  chloride.  It  re- 
duces the  former,  first  to  mercurous  chloride,  and  it  then 
abstracts  the  chlorine  from  this,  leaving  metallic  mercury  : 

2HgCl2  +  SnCl2  =  2HgCl  +  SnCl4  ; 
2HgCl  +  SnCl2  =  2Hg      +  SnCl4. 

Stannous  chloride  is  an  excellent  mordant,  and  is  ex- 
tensively used  by  the  dyers.  It  unites  with  other  chlo- 
rides, forming  double  chlorides  of  the  general  formula 
SnCl2.2MCl  or  M2SnCl4,  which  are  analogous  to  the  salts 
of  an  unknown  stannous  acid,  H2SnO2. 


642  INORGANIC  CHEMISTRY. 

Stannic  Chloride,  SnCl4,  formed  by  treating  tin  with 
chlorine,  is  a  colorless  liquid  which  boils  at  114°,  and 
the  specific  gravity  of  its  vapor  is  that  required  by  a 
compound  of  the  formula  SnCl4.  In  the  air  it  gives  off 
fumes  in  consequence  of  the  action  of  moisture.  It  has 
long  been  known  by  the  name  spiritus  fumans  Libavii, 
which  has  reference  to  its  fuming  quality  and  to  the  fact 
that  it  was  discovered  by  Libavius.  It  combines  with- 
water,  forming  a  number  of  crystallized  hydrates.  When 
its  solution  in  water  is  boiled  stannic  acid  is  precipitated  : 

SnCl4  +  3H2O  =  H2SnO3  +  4HC1. 

Probably  the  normal  acid,  Sn(OH)4,  is  first  formed,  and 
this  then  breaks  down  with  loss  of  water  to  form  the 
ordinary  acid : 

SnCl4  +  4H2O  =  Sn(OH)4  +  4HC1 ; 
Sn(OH)4  ==•  SnO(OH)2  +  H2O. 

Stannic  chloride  is  used  as  a  mordant.  With  other 
chlorides  it  forms  the  chlorostannates  of  the  general  for- 
mula M2SnCl6  or  SnCl4.2MCl,  of  which  the  ammonium 
salt,  (NH4)2SnCl6,  or  pink  salt,  is  the  best  known.  This 
is  manufactured  for  use  in  making  cotton -prints.  Similar 
fluorine  compounds,  HIQ  fluostannates,  are  also  easily  ob- 
tained, as  K2SnF6,  Na2SnF6,  etc. 

Stannous  Hydroxide,  Sn(OH)2,  is  not  known.  When  a 
solution  of  stannous  chloride  is  treated  with  potassium 
carbonate  a  precipitate  of  the  composition  H2Sn2O3  is 
formed,  which  is  a  derivative  of  the  hydroxide  Sn(OH)2, 
as  is  shown  by  the  formula  HO-Sn-O-Sn-OH,  which 
probably  expresses  its  constitution.  It  readily  loses 
water  and  passes  over  into  stannous  oxide,  SnO,  which  is 
a  black  powder. 

Stannic  Hydroxide,  Sn(OH)4,  is  perhaps  formed  when  a 
solution  of  stannic  chloride  in  water  is  boiled.  The  pre- 
cipitate obtained  has,  however,  the  composition  H2SnO3, 
and  this  is  known  as  stannic  acid.  Stannic  acid  is  pre- 
cipitated also  by  treating  a  solution  of  a  stannate  with 


METASTANNIC  ACID-STANNOUS  OXIDE.  643 

just  enough  of  an  acid  to  effect  decomposition.  The  de- 
composition with  hydrochloric  acid  takes  place  as  repre- 
sented in  the  equation 

Na2SnO3  +  2HC1  =  2NaCl  +  H2SnO3. 

The  compound  thus  obtained  is  insoluble  in  water,  but 
is  easily  soluble  in  hydrochloric,  nitric,  and  sulphuric 
acids,  and  in  the  caustic  alkalies.  With  the  alkalies  it 
forms  stannates,  as  sodium  stannate,  Na2SnO3,  and  potas- 
sium stannate,  K2SnO3.  The  former  is  made  on  the  large 
scale,  and  is  known  as  preparing  salt. 

Metastannic  Acid. — When  tin  is  treated  with  concen- 
trated nitric  acid  it  is  converted  into  a  white  powder  which 
is  insoluble  in  water  and  in  acids,  but  nevertheless  seems 
to  be  a  hydroxide  of  tin  of  the  same  composition  as  stannic 
acid.  This  is  known  as  metastannic  acid.  With  alkalies  it 
forms  salts  which  in  properties  and  composition  are  en- 
tirely different  from  the  stannates.  They  are  known  as 
the  metastannatcs.  Two  sodium  salts  are  known,  which 
differ  in  composition.  One  of  these  has  the  composition 
Na2Sn6On,  the  other  is  Na2Sn9O19.  They  are  probably 
derived  from  acids  analogous  to  the  polysilicic  acids, 
which  may  be  called  poly  stannic  acids.  The  question  as 
to  the  composition  of  metastannic  acid  is  an  open  one. 
When  heated  it  is  converted  into  the  oxide  SnO2.  When 
treated  with  concentrated  hydrochloric  acid  it  is  con- 
verted into  a  compound  containing  chlorine  which  is 
soluble  in  water,  and  the  solution  contains  stannic  chlo- 
ride. When  this  solution  is  treated  with  sulphuric  acid 
stannic  sulphate  is  thrown  down,  and  when  the  solution 
is  boiled  this  salt  is  decomposed,  leaving  stannic  acid. 

Stannous  Oxide,  SnO,  is  formed  from  the  correspond- 
ing hydroxide  when  this  is  heated  in  a  current  of  carbon 
dioxide.  When  a  solution  of  stannous  chloride  is  treated 
with  caustic  soda  stannous  hydroxide  is  first  precipitated, 
and  this  dissolves  in  an  excess  of  the  caustic  soda.  When 
the  solution  is  boiled  the  salt  contained  in  it  is  decom- 
posed, and  black  stannous  oxide  is  thrown  down  in  crys- 
talline form.  When  heated  in  the  air  stannous  oxide 
takes  fire,  and  is  converted  into  stannic  oxide. 


644  INORGANIC  CHEMISTEY. 

Stannic  Oxide,  SnO2,  as  has  been  stated,  is  the  princi- 
pal form  in  which  tin  occurs  in  nature.  The  mineral  is 
known  as  cassiterite  or  tin-stone.  It  is  found  at  Corn- 
wall, England  ;  on  the  East  Indian  islands  Banca  and 
Biliton  ;  and  in  Malacca.  It  crystallizes  in  the  tetragonal 
system,  and  is  generally  colored  from  brown  to  black. 
It  is  formed  as  a  white  powder  by  burning  tin  in  the  air, 
and  by  heating  the  different  varieties  of  stannic  hydrox- 
ide. It  is  infusible,  and  is  not  acted  upon  by  concen- 
trated hydrochloric  or  nitric  acid.  Concentrated  sul- 
phuric acid,  however,  converts  it  into  a  gelatinous  liquid 
from  which  stannic  oxide  is  precipitated  by  water. 

Stannous  Sulphide,  SnS,  which  is  formed  when  hydro- 
gen sulphide  is  passed  into  a  solution  of  stannous  chlo- 
ride, is  a  brownish-black  powder.  When  treated  with 
the  soluble  sulphides  and  sulphur,  or  with  a  soluble 
polysulphide,  it  dissolves,  forming  a  sulphostannate  as 
represented  in  these  equations  : 

SnS  +  (NH4)2S  +  S  -  (NH4)2SnS3 ; 
SnS  +  (NH4)2S2  =  (NH4)2SnS3. 

Stannic  Sulphide,  SnS2. — This  compound  is  obtained 
in  crystallized  form  by  heating  together  tin-filings,  sul- 
phur, and  dry  ammonium  chloride  ;  and  in  amorphous 
form  by  treating  a  solution  of  stannic  chloride  with  hy- 
drogen sulphide.  In  the  former  case  it  is  a  golden- 
yellow  crystalline  substance  ;  in  the  latter  a  yellow 
powder.  The  crystalline  variety  is  known  as  mosaic  gold. 
"When  heated  to  a  high  temperature  it  is  converted  into 
stannous  sulphide.  The  precipitated  variety  is  easily 
dissolved  by  concentrated  hydrochloric  acid,  and  con- 
verted into  metastannic  acid  by  concentrated  nitric  acid. 
The  crystallized  variety  is  not  soluble  in  hydrochloric 
acid,  and  is  affected  but  slightly  by  nitric  acid.  The 
crystallized  variety,  or  mosaic  gold,  is  used  as  a  pigment, 
particularly  for  bronzing.  Stannic  sulphide  dissolves 
easily  in  the  soluble  sulphides,  forming  sulphostan- 
nates : 

SnS,  +  (NH,),S  =  (NH,),SnS, ; 

SnS,  +  K,S         =  K.SnS,. 


SULPHOSTANNATES,  ETC.— REACTIONS.  645 

The  sulphostannates  are  perfectly  analogous  in  com- 
position to  the  stannates,  differing  from  them  simply  by 
containing  three  atoms  of  sulphur  in  place  of  the  three 
atoms  of  oxygen.  Although  the  comparison  has  been 
repeatedly  made  before  in  this  book,  it  is  perhaps  not 
superfluous  again  to  call  attention  to  the  analogy  between 
the  stannates,  sulphostannates,  chlorostannates,  and 
fluostannates,  all  of  which  are  easily  obtained,  and  are 
well-characterized  compounds.  The  formulas  of  some 
representatives  of  the  four  classes  are  here  placed  side 
by  side : 

K2Sn03  K2SnS3  K2SnCl6  K2SnF6 

Na2SnO3  Na2SnS3  Na2SnCl6  Na2SnF6 

When  a  solution  of  a  sulphostannate  is  treated  with 
an  acid  a  yellow  precipitate  is  formed.  This  becomes 
darker,  and  if  filtered  off  and  dried  it  forms  a  gray  mass. 
It  is  thought  that  this  may  be  sulphostannic  acid, 
H2SnS3.  It  breaks  down  readily  into  hydrogen  sulphide 
and  stannic  sulphide. 

Stannous  and  Stannic  Salts. — With  acids,  tin  forms  a 
few  salts,  but  they  are  not  particularly  well  characterized, 
and  the  stannic  salts  are  easily  decomposed  by  water. 
Stannous  sulphate,  SnSO4,  is  formed  by  dissolving  tin  in 
warm  concentrated  sulphuric  acid.  Stannic  sulphate, 
Sn(SO4)2,  is  precipitated  when  sulphuric  acid  is  added 
to  a  dilute  solution  of  stannic  chloride.  It  is  decomposed 
by  hot  water,  forming  stannic  acid. 

Reactions  which  are  of  Special  Value  in  Chemical  Analy- 
sis.— The  reactions  with  hydrogen  sulphide,  and  the  con- 
duct of  the  precipitated  sulphides  when  treated  with 
ammonium  sulphide  or  polysulphide,  are  constantly 
utilized  for  analytical  purposes  when  tin  is  present. 

The  reducing  action  of  stannous  compounds  serves  to 
distinguish  them  from  stannic  compounds. 

The  conduct  of  solutions  of  stannous  and  stannic  com- 
pounds towards  caustic  alkalies  has  been  explained  above. 

The  carbonates  of  the  alkali  metals  precipitate  the  hy- 
droxides, and  these  do  not  dissolve  in  an  excess  of  the 
carbonate.  Metallic  zinc  precipitates  the  tin  from  a  so- 


G46  INORGANIC  CHEMISTRY. 

lution  of  a  tin  compound  as  a  spongy  mass.  If  the  pre- 
cipitation is  carried  on  in  a  platinum  vessel  the  plati- 
num is  not  colored  by  it.  By  this  means  tin  can  be  dis- 
tinguished from  antimony,  which  is  reduced  under  the 
same  circumstances,  but  is  deposited  as  a  black  coating 
upon  the  platinum. 

LEAD,  Pb  (At.  Wt.  205.36). 

General. — The  basic  properties  of  lead  are  stronger, 
and  its  acidic  properties  weaker,  than  those  of  tin.  Its 
principal  compounds  are  those  in  which  it  acts  as  a  base- 
forming  element.  The  compounds  in  which  it  is  quad- 
rivalent, such  as  PbO2,  are  comparatively  unstable,  and 
when  treated  with  acids  they  readily  pass  over  into  the 
compounds  of  the  series  in  which  the  lead  is  bivalent. 
Thus  lead  oxide  itself  readily  gives  up  oxygen  when 
treated  with  acids,  and  yields  salts  of  bivalent  lead. 

Forms  in  which  Lead  occurs  in  Nature. — Lead  occurs  in 
nature  as  the  sulphide,  PbS,  which  is  known  as  galenite 
or  galena.  Other  natural  compounds  of  the  metal  are 
the  carbonate,  PbCO3,  known  as  cerussite ;  the  phos- 
phate, Pb3(PO4)2 ;  the  chromate,  PbCrO4,  or  crocoisite  ; 
and  the  molybdate,  PbMoO4,  or  wulfenite. 

Metallurgy. — Most  of  the  lead  in  the  market  is  ob- 
tained from  the  sulphide,  and  as  most  of  the  sulphide 
contains  silver  both  metals  have  to  be  considered  in  the 
treatment  of  the  ore.  Under  the  head  of  Silver  (which 
see)  reference  was  made  to  the  methods '  by  which  this 
metal  was  separated  from  the  lead  after  both  have  been 
separated  from  their  compounds.  Here  it  will  only  be 
necessary  to  show  how  the  metals  are  extracted  together 
from  the  ore.  This  is  accomplished  in  one  of  two  ways  : 

(1)  By  heating  the  sulphide  with  iron,  when  the  latter 
combines  with  the  sulphur,  forming  iron  sulphide,  while 
the  lead  is  set  free. 

(2)  By  roasting   the  sulphide  until  it  is  partly  con- 
verted into  lead  oxide  and  lead  sulphate,  and  then  heat- 
ing the  mixture  without  access  of  air,  when  two  reactions 
take  place,  which  are  represented  in  these  equations : 

PbS  +  2PbO  =  3Pb  +  SO2; 
PbS  +  PbS04  =  2Pb  +  2SOa. 


PROPERTIES  OF  LEAD.  647 

The  lead  is  thus  set  free,  and  the  sulphur  is  driven  off  as 
sulphur  dioxide. 

Properties. — Lead  is  a  bluish-gray  metal,  with  a  high 
lustre.  It  is  very  soft-,  and  not  very  strong ;  melts  at 
325°,  and  has  the  specific  gravity  11.37.  At  a  high  tem- 
perature it  is  converted  into  vapor.  When  heated  in 
contact  with  the  air  it  becomes  covered  with  a  layer  of 
oxide,  as  can  be  seen  in  a  vessel  containing  the  molten 
metal.  In  the  air,  at  ordinary  temperatures,  it  is  tar- 
nished in  consequence  of  the  formation  of  a  suboxide  of 
the  composition  Pb2O.  The  formation  of  this  compound 
can  be  observed  very  readily  by  cutting  a  piece  of  lead 
with  a  knife.  At  first  the  freshly- cut  surface  has  a  high 
lustre,  but  this  soon  grows  dim  and  acquires  a  bluish 
color.  Pure  water  acts  upon  lead  when  air  has  access 
to  it,  and  some  of  the  lead  dissolves.  If  the  water  con- 
tains salts  in  solution,  such  as  calcium  carbonate,  gyp- 
sum, etc.,  or  if  it  contains  carbon  dioxide,  it  acts  only 
very  slightly  upon  the  metal ;  further,  water  which  con- 
tains organic  matter  in  a  state  of  decomposition  dis- 
solves lead  with  comparative  ease.  Nitric  acid  dissolves 
Jead  readily;  as,  however,  lead  nitrate  is  insoluble  in 
nitric  acid,  it  is  necessary  to  use  comparatively  dilute 
acid.  Concentrated  sulphuric  acid  dissolves  lead  to 
some  extent,  and  therefore  a  little  lead  sulphate  is  some- 
times contained  in  commercial  sulphuric  acid.  When 
such  a  solution  is  diluted  with  water  the  sulphate  is  pre- 
cipitated. Hence,  whenever  the  commercial  acid  is  di- 
luted with  water  it  becomes  turbid,  and  after  standing 
for  a  time  it  becomes  clear,  as  the  lead  sulphate  settles. 
Hydrochloric  acid  acts  only  slightly  upon  lead.  Acetic 
acid  dissolves  the  metal  very  readily.  It  is  precipitated 
in  metallic  form  from  a  solution  of  one  of  its  salts  by 
metallic  zinc.  The  formation  is  sometimes  very  beauti- 
ful, especially  if  the  zinc  is  suspended  in  the  solution. 
It  is  called  the  "  lead  tree,"  or  Arbor  Saturni.  The 
action  consists  in  a  replacement  of  the  lead  by  the  zinc. 
After  the  action  is  complete  all  the  lead  is  deposited  as 
metallic  lead,  and  the  zinc  has  entered  into  its  place, 


648  INORGANIC  CHEMISTRY. 

forming  a  salt  which  remains  in  solution.  Thus,  if  lead 
nitrate  is  used,  zinc  nitrate  is  in  the  solution. 

Applications. — Lead  is  extensively  used  for  a  variety 
of  purposes,  as,  for  example,  for  making  sulphuric-acid 
chambers,  for  evaporating-pans  for  alum  and  sulphuric 
acid,  for  shot,  for  water-pipes,  and  in  making  alloys. 
The  use  of  lead  water-pipes  is  a  matter  of  much  im- 
portance from  the  sanitary  point  of  view,  as  is  evident 
from  the  statements  above  made  concerning  the  action  of 
water  upon  the  metal.  Ordinary  drinking-water  acts  under 
most  circumstances  only  very  slightly  upon  lead,  and  not 
enough  is  dissolved  to  be  dangerous  to  those  using  the 
water.  At  the  same  time  circumstances  may  at  any 
time  arise  which  will  increase  the  solvent  power  of  the 
water,  and  thus  cause  serious  results ;  and  it  would  un- 
doubtedly be  better  if  the  use  of  such  pipes  could  be 
entirely  avoided  in  cases  in  which  the  water  is  to  be 
used  for  drinking  purposes. 

Lead  Chloride,  PbCl2,  is  formed  when  hydrochloric 
acid  or  a  soluble  chloride  is  added  to  a  cold  solution  of 
a  lead  salt,  and  appears  as  a  white  precipitate.  It  is 
soluble  in  hot  water,  and  is  deposited  in  the  form  of  long, 
needle-shaped  crystals  when  the  solution  cools.  It  occurs 
in  nature  in  small  quantity,  as  cotunnite. 

Lead  Tetrachloride,  PbCl4,  is  a  liquid  that  forms  a 
crystalline  mass  at  —  15°.  It  is  decomposed  into  the 
dichloride  and  chlorine  at  the  ordinary  temperature. 
At  105°  this  takes  place  with  explosion.  It  is  formed  by 
passing  chlorine  into  concentrated  hydrochloric  acid 
containing  lead  chloride.  After  saturation,  ammonium 
chloride  is  added,  when  the  double  salt,  PbCl4.  2NH4C1, 
separates  in  yellow  crystals.  Ice-cold  concentrated  sul- 
phuric acid  decomposes  it  into  ammonium  sulphate, 
hydrochloric  acid,  and  lead  tetrachloride,  which  collects 
as  a  yellow  oil  below  the  sulphuric  acid. 

Lead  Iodide,  PbI2,  is  a  yellow  substance  which  crystal- 
lizes from  water  in  beautiful  lustrous  laminae.  It  is  pre- 
cipitated when  potassium  iodide  is  added  to  a  solution 
of  a  lead  salt.  It  dissolves  in  potassium  iodide  and 
forms  a  salt  of  the  formula  PbI2.KI  or  KPbI3.  There 


OXIDES  OF  LEAD.  649 

is  also  a  hydrogen  compound  of  the  formula  H2PbI4, 
that  is  formed  by  dissolving  lead  iodide  in  hydriodic 
acid. 

Lead  Hydroxide,  Pb(OH)2,  is  not  known,  but  when  a  lead 
salt  is  treated  with  a  soluble  hydroxide  a  compound  close- 
ly related  to  this  hydroxide  is  formed.  This  varies  some- 
what in  composition,  according  to  the  method  by  which 
it  is  made  ;  but  it  is  usually  either  HO-Pb-O-Pb-OH  or 
HO-Pb-O-Pb-O-Pb-OH. 

Oxides  of  Lead. — Lead  forms  four  distinct  compounds 
with  oxygen,  the  formulas  and  names  of  which  are  as 
follows :  lead  suboxide,  Pb2O  ;  lead  oxide,  PbO ;  lead  ses- 
quioxide,  Pb2O3 ;  and  lead  peroxide,  PbO2. 

Lead  Suboxide,  Pb2O. — This  compound,  as  has  been 
stated,  is  formed  when  lead  is  exposed  to  the  air.  It  has 
been  made  in  pure  condition,  and  is  a  black  powder. 
When  treated  with  acids  it  yields  salts,  and  lead  sepa- 
rates. Thus  with  hydrochloric  acid  lead  chloride  is 
formed,  as  represented  in  the  equation 

Pb20  +  2HC1  =  PbCl2  +  H2O  +  Pb. 

Lead  Oxide,  PbO. — This  compound  is  formed  by  heat- 
ing lead  nitrate,  and  is  then  left  behind  in  the  form  of  a 
yellow  powder.  If  heated  to  melting  it  solidifies,  form- 
ing a  yellowish  or  reddish  mass  known  as  litharge.  This 
is  formed  in  large  quantity  in  the  process  of  separating 
silver  from  lead.  It  will  be  remembered  that,  in  order 
to  remove  the  last  portions  of  lead  from  silver,  the  alloy 
is  melted  and  air  blown  upon  it  in  this  condition.  Un- 
der these  circumstances,  the  lead  is  converted  into  the 
oxide  while  the  silver  remains  unchanged.  The  lead 
oxide  thus  formed  is  the  litharge  found  in  the  market. 
The  oxide  can  be  obtained  in  crystallized  form  by  heat- 
ing a  solution  of  the  ordinary  oxide  in  dilute  caustic  soda 
or  caustic  potash ;  and  by  boiling  lead  hydroxide  with  a 
quantity  of  caustic  alkali  insufficient  to  dissolve  it.  In 
the  powdered  condition  lead  oxide  attracts  carbon  diox- 
ide from  the  air.  With  acids  it  forms  salts  which  in 
some  respects  resemble  those  of  barium  and  strontium. 
With  the  strongest  bases  it  forms  salts  similar  to  those 


650  INORGANIC  CHEMISTRY. 

formed  by  zinc.  This  is  seen  in  the  solubility  of  the 
hydroxide  in  sodium  and  potassium  hydroxide,  which  is 
due  to  the  formation  of  compounds  known  as  plumbites  : 

Pb(OH)2  +  2KOH  =  Pb(OK)2  +  2H2O. 

Heated  with  silicon  dioxide  it  forms  a  silicate  which  is 
easily  fusible.  Lead  oxide  is  used  extensively  in  the 
manufacture  of  flint  glass  for  optical  purposes,  as  was 
described  under  Glass  (which  see).  It  also  finds  appli- 
cation in  glass  painting  and  porcelain  painting,  and  is 
used  in  the  manufacture  of  lead  salts,  particularly  "  sugar 
of  lead,"  which  is  an  acetate. 

Lead  Sesquioxide,  Pb2O3,  is  obtained  by  bringing  to- 
gether lead  acetate  and  caustic  soda,  and  treating  the 
solution  with  sodium  hypochlorite  ;  and  is  thrown  down 
as  a  reddish-yellow  powder.  The  constitution  of  the 
compound  is  unknown.  It  has  been  suggested  that  it 
may  be  a  lead  salt  of  plumbic  acid,  as  represented  in  the 

formula    PbO<Q>Pb,  plumbic  acid  being,  as  will  be 

seen, 

Lead  Peroxide,  PbO2,  is  formed  by  treating  minium  or 
red  lead  with  dilute  nitric  acid.  Minium  has  the  compo- 
sition, Pb3O4.  When  treated  with  nitric  acid,  a  part  dis- 
solves as  lead  nitrate,  and  lead  peroxide  remains  behind, 
as  represented  in  the  equation  : 

Pb3O4  +  4HNO3  =  PbOa  +  2Pb(N03)2  +  2H2O. 

The  peroxide  is  formed  in  general  by  the  action  of  oxidiz- 
ing agents  upon  the  lower  oxides  of  lead.  One  of  the 
most  convenient  methods  for  making  it  consists  in  treat- 
ing lead  acetate  with  a  filtered  solution  of  bleaching- 
powder.  It  is  a  dark-brown  powder,  insoluble  in  water. 
When  ignited  it  loses  half  of  its  oxygen,  and  it  gives  up 
its  oxygen  readily  to  other  substances.  Towards  hydro- 
chloric acid  it  acts  like  manganese  dioxide,  giving  lead 
chloride  and  chlorine  according  to  the  equation 

PbOa  +  4HC1  =  PbCl2  +  2H20  +  C12. 


RED  LEAD.  651 

It  appears  probable  that  the  tetrachloride  is  first 
formed,  and  that  this  then  breaks  down  into  the  dichlo- 
ride  and  chlorine.  When  the  peroxide  is  treated  in  the 
cold  with  hydrochloric  acid  it  dissolves,  and  when  this 
solution  is  heated  it  gives  off  chlorine.  Further,  when  it 
is  treated  with  caustic  alkalies  lead  peroxide  is  thrown 
down. 

Lead  peroxide  dissolves  in  concentrated  caustic  potash 

OTC 

and  forms  a  salt  of  the  formula  K9PbO3,  or  PbO<oK' 

analogous  to  potassium  stannate,  K2SnO3,  silicate, 
K2SiO3,  carbonate,  K2CO3,  etc.  Other  salts  derived  from 
the  acid  PbO(OH)2  are  known,  and  are  called  plumbates. 
Red  Lead,  Minium. — When  lead  oxide  is  heated  gently 
in  the  air  it  takes  up  oxygen,  and  is  converted  into  the 
red  compound  known  as  minium  or  red  lead.  The  com- 
mercial article  of  this  name  varies  in  composition,  but 
approximates  to  that  represented  by  the  formula 
Pb3O4,  and,  if  the  oxide  is  slowly  heated,  the  amount  of 
oxygen  taken  up  is  that  required  to  form  a  compound  of 
the  above  formula.  Red  lead  varies  in  color  from  red 
to  yellowish,  according  to  the  method  of  preparation. 
When  heated,  it  becomes  dark,  but  the  red  color  appears 
again  on  cooling.  When  heated  to  a  high  temperature, 
it  loses  oxygen  and  yields  lead  oxide  : 

Pb,04  =  3PbO  +  O. 

When  treated  with  dilute  nitric  acid,  lead  nitrate  is 
formed,  and  lead  peroxide  is  left  undissolved.  As  re- 
gards the  relation  existing  between  minium  and  the 
other  oxides  of  lead,  no  positive  statement  can  be  made. 
The  evidence  points  to  the  conclusion  that  it  is  a  chemi- 
cal compound  and  not  a  mixture.  Dilute  acetic  acid 
does  not  dissolve  it,  while  this  acid  does  dissolve  the 
monoxide.  It  has  been  suggested  that  it  is  a  lead  salt 
of  normal  plumbic  acid,  Pb(OH)4,  as  represented  in  the 


formula 


'O 
O 


>Pb 


o>pb 


As  partial  experimental  evidence 


652  INORGANIC  CHEMISTRY. 

in  support  of  this  view,  the  fact  may  be  mentioned  that  a 
compound  similar  to  red  lead  is  formed,  when  a  solu- 
tion of  potassium  plumbate  is  treated  with  a  solution 
of  lead  oxide  in  potassium  hydroxide.  In  solution,  the 
potassium  salt  probably  has  the  constitution  repre- 

roH 

OH 

sented  by  the  formula  Pb  -j  Q-^-.      When  this  is  treated 


OK 

with  lead  oxide  the  corresponding  lead  salt  should  be 
formed.  —  Red  lead  is  used  as  a  pigment,  and  sometimes 
in  place  of  litharge  when  an  oxide  of  lead  is  needed:  as 
in  the  manufacture  of  glass,  as  a  flux  in  the  manufacture 
of  porcelain,  etc. 

Lead  Sulphide,  PbS  —  This  has  already  been  referred 
to  as  the  principal  compound  from  which  lead  is  ob- 
tained. The  natural  variety  is  called  galena  or  galenite. 
It  is  formed  in  the  laboratory  as  a  black  precipitate,  when 
hydrogen  sulphide  is  passed  into  a  solution  of  a  lead  salt. 
When  heated  in  the  air,  as  in  the  roasting  of  galenite, 
the  sulphur  passes  off  as  sulphur  dioxide,  and  the  lead 
is  converted  into  oxide.  Concentrated  hydrochloric  acid 
dissolves  it.  Concentrated  nitric  acid  converts  it  into  the 
sulphate.  When  hydrogen  sulphide  is  conducted  into  a 
weak  acid  solution  of  lead  chloride,  a  compound  contain- 
ing lead,  sulphur,  and  chlorine  is  precipitated,  the  com- 
position of  which  is  approximately  that  represented  by 
the  formula  3PbS.PbCl2,  and  this  has  a  red  or  a  yellow 
color,  according  to  the  conditions. 

Lead  Nitrate,  Pb(NO3)2.  —  The  nitrate  is  easily  made 
by  dissolving  lead,  lead  oxide,  or  carbonate  in  nitric  acid. 
The  salt  crystallizes  well,  and  is  easily  soluble  in  water. 
It  is  difficultly  soluble  in  dilute  nitric  acid,  and  insoluble 
in  concentrated  nitric  acid,  resembling  in  this  respect 
barium  nitrate.  It  is  decomposed  by  heat,  giving  nitro- 
gen peroxide,  NO2,  and  lead  oxide. 

Lead  Carbonate,  PbCO3.  —  The  carbonate  occurs  in  na- 
ture as  cerussite,  crystallized  in  forms  which  are  the 
same  as  those  of  barium  carbonate,  and  of  that  variety  of 
calcium  carbonate  known  as  aragonite.  It  can  be  ob- 


LEAD  CARBONATE.  653 

tained  by  adding  lead  nitrate  to  a  solution  of  ammonium 
carbonate,  but,  when  solutions  of  lead  salts  are  treated 
with  the  secondary  carbonates  of  the  alkali  metals,  pre- 
cipitates of  basic  carbonates  are  always  obtained.  When 
an  excess  of  sodium  carbonate  is  added  to  a  solution 
of  lead  nitrate,  the  precipitate  has  the  composition 
HO-Pb-0-CO-O-Pb-O-CO-O-Pb-OH,  or  3PbO.2CO2 
+  H2O.  The  salts  usually  obtained  are  more  complicated 
than  this,  but  the  relations  between  them  and  lead  oxide 
and  carbonic  acid  are  of  the  same  kind.  Basic  lead  car- 
bonate is  prepared  and  used  extensively,  under  the  name 
of  white  lead,  as  a  pigment.  It  is  manufactured  by  differ- 
ent methods.  The  principal  ones  are  the  following : 

(1)  The  Dutch  Method. — This  consists  in  exposing  sheets 
of  lead  wound  in  spirals  to  the  action  of  vinegar,  air, 
and  carbon  dioxide  from  decaying  organic  matter.     The 
spirals  of  sheet  lead  are  placed  in  earthenware  vessels, 
on  the  bottom  of  which,  but  not  in  contact  with  the  lead, 
the  vinegar  is  placed.     The  vessels  thus  arranged  are 
placed  in  beds  of  horse  manure.     In  consequence  of  de- 
composition, which  is  set  up  in  the  manure,  carbon  diox- 
ide is  given  off  slowly,  and  enough  heat  is  generated  to 
start  the  action  upon  the  lead.     The  chemical  changes 
involved  in  the  process  are,  mainly,  the  formation  of  a 
basic  acetate  of  lead,  and  the  subsequent  decomposition 
of  this  by  carbon  dioxide,  forming  a  basic  carbonate, 
and  leaving  the  acetic  acid  free  to  act  upon  a  further 
quantity  of  lead. 

(2)  The  French  Method. — In  this  method  a  solution  of 
basic  lead  acetate  is  prepared  by  treating  a  solution  of 
the  neutral  salt  with  lead  oxide.     This  is  then  decom- 
posed by  passing  carbon  dioxide  into  it,  when  a  basic 
carbonate  is  thrown  down.     The  carbon  dioxide  is  gen- 
erally made  by  burning  coke. 

(3)  The  English  Method. — This  is  a  modification  of  the 
Dutch  method,  and  differs  from  it  chiefly  in  the  replace- 
ment of  manure  by  spent  tan  in  a  state  of  fermentation, 
and  the  use  of  dilute  acetic  acid  in  place  of  vinegar. 
There  is  less  risk  of  discoloration   in  consequence  of 
the  formation  of  sulphuretted  hydrogen,  but  the  fermen- 


654  INORGANIC  CHEMISTRY. 

tation  takes  place  more  slowly,  and  the  whole  process, 
therefore,  requires  a  longer  time. 

The  composition  of  white  lead  is  not  always  the 
same.  That  prepared  by  precipitating  a  solution  of  basic 
lead  acetate  with  carbon  dioxide  has  the  composition 
Pb(OH)2.3PbCO3  ;  and  that  prepared  by  the  Dutch 
method  has  the  composition  Pb(OH)2.2PbCO8  ;  or  these 
may  be  expressed  structurally  by  the  formulas 

Pb<o 
t-'» 

and 


Pb<OH  OH 

An  objection  to  white-lead  paint  is  that  it  turns  dark 
under  the  influence  of  hydrogen  sulphide.  It  also  turns 
yellow  in  consequence  of  the  action  of  some  substance 
contained  in  the  oil  with  which  the  lead  carbonate  is 
mixed. 

Lead  Sulphate,  PbSO4,  occurs  to  some  extent  in  nature. 
It  is  formed  by  adding  sulphuric  acid  or  a  soluble  sul- 
phate to  a  solution  of  a  lead  salt,  and  by  oxidation  of 
lead  sulphide.  Like  barium  sulphate,  it  is  practically 
insoluble  in  water.  As  stated  above,  it  is  somewhat 
soluble  in  concentrated  sulphuric  acid,  and  it  is  there- 
fore always  found  in  the  concentrated  acid  of  commerce. 
Nitric  acid  and  hydrochloric  acid  dissolve  it  in  consider- 
able quantity.  It  dissolves  further  quite  readily  in  solu- 
tions of  some  ammonium  salts,  as  in  ammonium  tartrate 
and  acetate.  When  heated  to  redness  it  is  partly  decom- 
posed with  loss  of  sulphur  trioxide. 

Reactions  which  are  of  Special  Value  in  Chemical  Analy- 
sis. —  The  reactions  of  lead  salts  with  the  soluble  hy- 
droxides, with  sulphuric  acid,  hydrochloric  acid,  hydro- 
gen sulphide,  soluble  carbonates,  potassium  chromate 
and  dichromate,  are  the  ones  which  are  principally  used 
in  analysis.  All  of  these  have  been  treated  of  in  this 
chapter,  with  the  exception  of  those  with  potassium 
chromate  and  dichromate,  which  will  be  taken  up  in  the 
chapter  on  Chromium  (which  see).  In  anticipation  it 


LANTHANUM— CERIUM.  655 

may  be  said  that  the  reactions  are  based  upon  the  fact 
that  lead  chromate,  PbCrO4,  like  barium  chromate,  is 
insoluble  in  water. 


The  elements  of  Family  V,  Group  A,  are  vanadium, 
columbium,  didymium,  and  tantalum.  As  they  are 
closely  related  to  the  members  of  Group  B,  of  the  same 
family,  they  were  treated  of  at  the  end  of  Chapter  XVIII. 
in  connection  with  the  members  of  the  phosphorus 
group.  Among  them  the  one  which  is  least  known  is 
didymium.  This  in  turn  is  more  or  less  closely  related 
to  two  other  elements  of  nearly  the  same  atomic  weight 
which  occur  in  Families  III  and  IV.  These  are  lantha- 
num and  cerium.  A  few  words  in  regard  to  these  three 
rare  elements  will  suffice  for  the  present  purpose. 

LANTHANUM,  CERIUM,  DIDYMIUM. 

These  three  elements  occur  together  in  several  rare 
minerals  of  Norway,  as  cerite,  gadolinite,  and  allanite. 
Cerite  is  a  silicate  of  the  three  metals,  and  its  composi- 

Ce.) 
tion  is  represented  by  the  formula  La4  V  (SiO,)3  +  3HaO. 

Di.  j 

It  is  probably  a  mixture  of  three  isomorphous  silicates. 
The  principal  constituent  is  cerium  silicate,  Ce4(SiO4)3. 
The  perfect  separation  of  the  constituents  of  the  mineral 
is  a  very  difficult  operation. 

Lanthanum,  La  (At.  Wt.  137.59),  forms  an  oxide  of  the 
formula  La2O3,  analogous  to  that  of  aluminium.  Its 
chloride  also  is  analogous  to  that  of  aluminium,  and  has 
the  composition  LaCls;  and  in  all  its  salts  it  acts  as  a 
trivalent  element. 

Cerium,  Ce  (At.  Wt.  139.1),  forms  two  series  of  com- 
pounds, in  one  of  which  it  is  trivaient,  resembling  lan- 
thanum and  the  other  members  of  the  aluminium  group  ; 
while  in  the  other  series  it  is  quadrivalent,  resembling 
silicon  and  the  other  members  of  the  silicon  group. 
The  formulas  of  some  of  the  principal  members  of  the 
first  series  are  as  follows : 

CeCl3,  Ce203,  and  Ce2(S04)3. 


656  INORGANIC  CHEMISTRY. 

Some  of  the  principal  members  of  the  second  series  are 
represented  by  the  formulas 

CeF4,  CeO2,  Ce(NO3)4,  and  Ce(SO4)a. 

Didymium,  Di  (At.  Wt.  142.1),  has  already  been  re- 
ferred to  on  page  351  in  connection  with  the  members 
of  Family  Y,  Group  A,  which  it  resembles  in  some  re- 
spects. In  most  of  its  compounds  it  is,  however,  triva- 
lent,  forming  compounds,  of  some  of  which  the  following 
are  the  formulas  : 

DiCls,  Di203,  Di(N03)3,  Di2(S04)3,  Di2(CO3)3,  etc. 

Praseodymium,  Pr,  and  Neodymium,  Nd. — While  the 
name  didymium  is  still  given  above,  and  this  substance 
dealt  with  as  though  it  were  an  element,  as  it  was  at  first 
held  to  be,  it  has  been  shown  by  Auer  von  Welsbach 
that  it  consists  of  two  very  similar  elements  to  which  he 
has  given  the  names  praseodymium  and  neodymium..  When 
the  double  nitrate  of  ammonium  and  didymium  is  re- 
peatedly recrystallized  it  is  separated  into  two  salts,  one 
of  which  is  green,  and  the  other  rose-colored.  When  the 
nitrate  or  oxalate  of  one  of  these  new  elements  is  ignited 
it  forms  a  black  oxide,  while  from  the  other  is  formed  an 
oxide  of  a  different  color.  The  element  that  gives  green 
salts  is  called  praseodymium,  and  the  other  neodymium. 
The  atomic  weights  of  these  elements  are  nearly  the 
same,  but  they  have  not  yet  been  accurately  determined. 


CHAPTER  XXXI. 

ELEMENTS  OF  FAMILY  VI,  GROUP  A : 
CHROMIUM— MOLYBDENUM— TUNGSTEN— URANIUM. 

General. — At  the  end  of  Chapter  XIY.,  in  connection 
with  the  elements  of  the  sulphur  group,  the  four  ele- 
ments which  form  the  subject  of  this  chapter  were  briefly 
referred  to,  for  the  reason  that  in  some  respects  they 
resemble  sulphur.  As  was  there  stated,  this  resem- 
blance "  is  seen  mainly  in  the  formation  of  acids  of  the 
formulas  H2CrO4,  H2MoO4,  H2WO4,  and  H2UO4 ;  and  the 
oxides  CrO3,  MoO3,  WO3,  and  UO3."  Further,  it  was 
stated  that  "  when  the  acids  of  chromium,  molybdenum, 
tungsten,  and  uranium  lose  oxygen,  they  form  com- 
pounds which  have  little  or  no  acid  character.  The  lower 
oxides  of  chromium  form  salts  with  acids,  and  these  bear 
a  general  resemblance  to  the  salts  of  aluminium,  iron, 
and  manganese.  The  chromates  lose  their  oxygen  quite 
readily  when  acids  are  present  with  which  the  chromium 
can  enter  into  combination  as  a  base-forming  element." 
"  Molybdenum  and  tungsten  do  not  form  salts  of  this 
character :  indeed  they  seem  to  be  practically  devoid  of 
the  power  to  form  bases.  Uranium,  on  the  other  hand, 
forms  some  curious  salts  which  differ  from  the  simple 
metallic  salts  which  we  commonly  have  to  deal  with. 
These  are  the  uranyl  salts  which  are  regarded  as  acids, 
in  which  the  hydrogen  is  either  wholly  or  partly  replaced 
by  the  complex  UO2,  which  is  bivalent.  Thus,  the  nitrate 
has  the  formula  UO2(NO3)2,  the  sulphate  (UO2)SO4,  etc. 
These  salts  are  derived  from  the  compound  UO2(OH)2, 
acting  as  a  base,  whereas  the  compound  has  also  dis- 
tinctly acid  properties."  That  member  of  the  group  the 
compounds  of  which  are  most  commonly  met  with  in 
the  laboratory  and  in  the  arts  is  chromium,  and  this  will 
receive  principal  attention  here. 

(657) 


658  INORGANIC  CHEMISTRY. 

CHROMIUM,  Or  (At.  Wt.  51.74). 

General. — This  element  forms  three  series  of  com- 
pounds, in  which  it  appears  to  be  respectively  bivalent, 
trivalent,  and  sexivalent.  Of  these  the  members  of  the 
series  in  which  it  is  trivalent  are  most  stable  under  ordi- 
nary circumstances.  Some  of  the  principal  members  of 
the  first  series,  or  the  chromous  compounds,  are  repre- 
sented by  the  formulas 

CrCl2,  Cr(OH)2,  CrSO4,  CrCO3. 

Of  the  second  series,  or  the  chromic  compounds,  some 
of  the  principal  members  are : 

CrCl3,  Cr203,  Cr2(SO4)3,  Cr(NO3)3,  KCr(SO4)2  +  12H2O. 

And,  finally,  the  members  of  the  third  series  are  derived 
from  the  oxide  CrO3,  and  they  are  for  the  most  part  salts 
of  the  acid  of  the  formula  H2CrO4,  known  as  chromic  acid, 
or  of  an  acid  of  the  formula  H2Cr2O7,  known  as  dichromic 
acid,  which  is  closely  related  to  chromic  acid. 

When  exposed  to  the  air  the  chromous  compounds 
are  converted  into  chromic  compounds,  and  they  are  in 
general  readily  converted  into  chromic  compounds  by 
the  action  of  oxidizing  agents,  as  cuprous  and  mercurous 
compounds  are  converted  into  cupric  and  mercuric  com- 
pounds. If  the  oxidation  takes  place  in  acid  solution 
the  limit  is  reached  when  a  chromic  salt  is  formed.  If, 
however,  the  action  takes  place  in  the  presence  of  a 
strong  base  the  limit  is  reached  in  the  formation  of  a 
chromate.  Thus,  suppose  chromous  oxide  to  be  treated 
with  an  oxidizing  agent  in  the  presence  of  sulphuric  acid, 
the  final  product  would  be  chromic  sulphate,  as  repre* 
sented  in  the  following  equations  : 

CrO  +  H2S04  =  CrS04  +  H2O  ; 

2CrS04  +  H2S04  +  O  =  Cr2(SO4)3  +  H2O. 

On  the  other  hand,  if  the  oxidation  takes  place  in  the 
presence  of  caustic  potash  the  final  product  is  potassium 
chromate,  as  shown  in  the  following  equation : 

CrO  +  2KOH  +  Oa  =  K2CrO4  +  H2O. 


CHROMIUM.  659 

When  a  chromate  is  treated  with  an  acid  it  tends  to 
pass  back  to  a  compound  of  the  chromic  series,  and  the 
change  involves  the  giving  up  of  oxygen.  Thus  when 
potassium  chromate  is  treated  with  sulphuric  acid  in  the 
presence  of  something  which  has  the  power  to  take  up 
oxygen,  potassium  and  chromium  sulphates  are  formed, 
and  oxygen  is  given  up,  thus  : 

2K2CrO4  +  5H2SO4  =  2K2SO4  -f  Cr2(SO4)3  +  5H2O  -f  3O. 

All  these  relations  will  be  more  fully  taken  up  in  the 
paragraphs  which  treat  of  the  individual  compounds. 

Forms  in  which  Chromium  Occurs  in  Nature. — The 
principal  form  in  which  chromium  occurs  in  nature  is 
the  mineral  chromite,  also  known  as  chromic  iron  and 
chrome  iron  ore.  This  has  the  composition  FeCr2O4, 
and,  as  will  be  pointed  out  below,  it  is  probably  analo- 
gous to  the  spinels  (see  p.  572),  being  an  iron  salt  of  the 
acid  CrO.OH,  which  may  be  called  metachromous  acid. 

CrO  O 
This  view  is   represented   by  the  formula   ^  Q*^>Fe. 

It  occurs  also  in  the  mineral  crocoisite,  which  is  lead 
chromate,  PbCrO4.  The  name  chromium  is  derived  from 
the  Greek  jpcSyua',  meaning  color ;  and  the  element  is  so 
called  because  most  of  its  compounds  are  colored. 

Preparation. — The  metal  is  obtained  by  the  electroly- 
sis of  chromic  chloride ;  by  decomposing  the  chloride 
by  means  of  sodium  in  the  form  of  vapor ;  and  by  treat- 
ing the  chloride  with  zinc. 

Properties. — Chromium  is  a  light-gray,  crystalline,  lus- 
trous, metallic-looking  substance ;  or  it  consists  of  mi- 
croscopic, lustrous  rhombohedrons  of  a  tin-white  color. 
It  is  very  hard,  and  difficultly  fusible.  When  heated  in 
the  air  it  is  oxidized  very  slowly,  but  in  the  flame  of  the 
oxyhydrogen  blowpipe  it  burns,  forming  chromic  oxide, 
Cr2O3.  It  is  easily  dissolved  by  hydrochloric  acid.  Cold 
sulphuric  acid  does  not  dissolve  it ;  the  hot  acid  does. 
Nitric  acid  does  not  affect  it.  When  treated  with  salts 
of  potassium  which  easily  give  up  their  oxygen,  as  the 
chlorate  and  nitrate,  it  is  converted  into  potassium 
chromate. 


660  INORGANIC  CHEMISTRY. 

Chromous  Chloride,  CrCl2,  is  formed  by  dissolving  the 
metal  in  hydrochloric  acid,  and  by  carefully  heating 
chromic  chloride  in  a  current  of  hydrogen.  It  forms 
white  crystals,  which  dissolve  in  water,  giving  a  blue 
solution.  This  solution  takes  up  oxygen  very  readily 
from  the  air,  and  the  compound  is  converted  into  others 
which  belong  to  the  chromic  series.  The  other  chro- 
mous  compounds  act  in  a  similar  way. 

Chromic  Chloride,  CrCl3. — This  compound  is  made  in 
solution  by  dissolving  chromic  hydroxide,  Cr(OH)3,  in 
hydrochloric  acid.  This  solution  has  a  dark-green  color. 
When  evaporated  to  a  sufficient  extent  crystals  of  the 
composition  CrCl3  +  6H2O  are  deposited.  If  these  are 
heated  in  the  air  they  undergo  decomposition  just  as 
aluminium  chloride  does,  and  the  product  left  behind  is 
chromic  oxide  : 

2CrCl3  +  3H2O  =  Cr2O3  +  6HC1. 

If,  however,  the  crystallized  chloride  is  heated  in  an 
atmosphere  of  chlorine  or  hydrochloric  acid,  the  water 
is  given  off,  and  the  anhydrous  chloride,  which  has  a 
beautiful  reddish  violet  color,  is  formed.  This  dissolves 
in  water  and  forms  a  green  solution.  But  if  the  dry 
chloride  thus  obtained  is  sublimed,  it  is  deposited  in 
lustrous  laminae  of  the  same  color  ;  and  this  variety 
is  insoluble  in  water  and  acids,  and  is  only  slowly 
acted  upon  by  boiling  alkalies.  This  insoluble,  crystal- 
lized variety  of  the  chloride  is  obtained  also  by  the  same 
method  as  that  used  in  making  aluminium  chloride,  that 
is,  by  passing  a  current  of  chlorine  over  a  heated  mix- 
ture of  carbon  and  chromic  oxide.  Although  it  is  called 
insoluble,  it  passes  gradually  into  solution  by  boiling 
with  water.  Further,  when  a  very  minute  quantity  of 
chromous  chloride  is  mixed  with  it,  it  dissolves  easily, 
and  forms  a  green  colored  solution. 

Chromic  chloride  unites  with  other  chlorides,  as  alu- 
minium chloride  does,  and  forms  double  chlorides, 
analogous  to  the  chlor-aluminates.  Examples  of  these 
are  the  compounds  of  the  formulas  CrCl3.KCl,  or 


CHROMIC  HYDROXIDE.  661 

KCrCl4;  CrCl3.2KCl,  or  K2CrCl6;  and  CrCl3.3KCl,  or 
K3CrCl6. 

Chromous  Hydroxide,  Cr(OH)2,  is  formed  as  a  brown- 
ish-yellow precipitate  by  adding  caustic  potash  to  a  solu- 
tion of  chromous  chloride.  It  easily  gives  up  hydrogen, 
and  is  converted  into  chromic  oxide  : 

2Cr(OH),  =  Cr.O.  +  H2O  +  H,.         |    , 

Chromic  Hydroxide,  Cr(OH)3. — When  ammonia  is  added 
to  a  solution  of  a  chromic  salt,  a  light-blue  voluminous 
precipitate,  which  has  the  composition  Cr(OH)3  -|-  2H2O, 
is  formed.  When  this  is  filtered  off  and  dried  in  a 
vacuum  it  loses  the  water  and  leaves  the  hydroxide. 
This  is  readily  converted  by  heat  into  a  compound  of  the 
formula  CrO.OH,  and  finally  into  chromic  oxide,  Cr2O3. 
The  green  precipitates  formed  in  solutions  of  chromic 
salts  by  sodium  and  potassium  hydroxides  always  con- 
tain some  of  the  alkali  metal  in  combination. 

Chromic  hydroxide,  like  aluminium  hydroxide,  dis- 
solves in  the  soluble  hydroxides,  and  forms  salts  known 
as  chromites,  which  are  derived  from  the  acid  CrO.OH. 
Thus  with  potassium  hydroxide  the  action  takes  place 
as  represented  in  the  equation 

/OH 

Crf-  OH  +  KOH  =  Cr^Xi^  +  2H2O. 
X)H 

If  the  solution  containing  potassium  or  sodium  chro- 
mite  is  boiled,  the  salt  is  decomposed  and  chromic  hy- 
droxide precipitated,  though  the  precipitate  thus  formed 
always  contains  some  of  the  alkali  metal  in  combination. 
It  will  be  noticed  that  in  this  respect  aluminium  and 
chromium  conduct  themselves  differently  towards  the 
alkaline  hydroxides. 

It  has  already  been  stated  that  chromite,  (CrO.O)2Fe, 
is  regarded  as  an  iron  salt  of  the  same  order  as  the  po- 
tassium salt  referred  to. 


662  INORGANIC  CHEMISTRY. 

Another  hydroxide  formed  by  heating  potassium  di- 
chromate  and  boric  .acid  together  has  the  composition 
represented  by  the  formula  Cr2O(OH)4  or  Cr4O3(OH)6. 
This  is  known  as  Guignet's  green.  The  relation  between 
the  normal  hydroxide  and  these  compounds  is  shown 
by  means  of  the  equations 

2Cr(OH)3  =  Cr2O(OH)4  +  H2O  ; 
4Cr(OH)3  =  Cr4O3(OH)6  +  3H2O. 

Chromic  Oxide,  Cr2Os,  is  formed  by  igniting  the  hy- 
droxides, and  is  most  readily  prepared  by  heating  a 
mixture  of  potassium  dichromate  and  sulphur.  The 
sulphur  is  oxidized,  and  with  the  potassium  forms 
potassium  sulphate,  while  the  chromic  acid  is  reduced 
to  the  form  of  the  oxide  Cr2O3.  It  can  be  obtained  in 
crystals.  As  ordinarily  obtained  it  is  a  green  powder, 
which  after  ignition  is  almost  insoluble  in  acids.  It  is 
dissolved,  however,  by  treatment  with  fusing  potassium 
sulphate.  The  oxide  colors  glass  green,  and  is  used  in 
painting  porcelain. 

Chromic  Sulphate,  Cr2(SO4)3,  is  made  by  dissolving  the 
hydroxide  in  concentrated  sulphuric  acid  when  it  is  de- 
posited in  purple  crystals  of  the  composition  Cr2(SO4)3  -f- 
15H2O.  If  the  solution  of  this  salt  is  boiled,  the  so- 
lution becomes  green,  and  crystals  cannot  be  obtained 
from  it.  But  by  standing  for  some  time  the  green  solu- 
tion becomes  reddish  purple  again,  and  yields  the  crys- 
tallized salt.  Other  salts  of  chromium  act  in  the  same 
way.  They  exist  in  two  varieties,  one  of  which  crystal- 
lizes and  is  reddish  purple  in  color,  while  the  other  does 
not  crystallize  and  is  green.  The  crystallized  salts  are 
converted  into  the  uncrystallized  green  salts  by  boiling, 
and  the  green  salts  are  converted  into  the  crystallized 
salts  by  standing. 

Chrome- Alums. — Chromic  sulphate,  like  aluminium  sul- 
phate, combines  with  other  sulphates,  such  as  potassium, 
sodium,  and  ammonium  sulphates,  and  forms  well-crys- 
tallized salts,  which  are  closely  analogous  to  ordinary 


CHROMIC  ACID  AND  THE  CHROMATES.  663 

alum.  They  all  contain  twelve  molecules  of  water,  as 
represented  in  -the  formulas  below  : 

Chrome-Alum, KCr(SO4)2        -f  12H0O 

Sodium  Chrome-Alum,  .  .  .  NaCr(SO4)2  +  12H2~O 
Ammonium  Chrome-Alum,  .  (NH4)Cr(SO4)2  +  12H2O 

The  potassium  compound  which  is  commonly  called 
chrome-alum  is  made  by  adding  a  reducing  agent,  such 
as  alcohol  or  sulphur  dioxide,  to  a  solution  of  potas- 
sium dichromate  containing  sulphuric  acid.  If  the  solu- 
tion is  heated  it  turns  green,  and  crystals  cannot  be  ob- 
tained from  it.  But  on  standing  for  a  considerable  time 
its  color  changes,  and  reddish-purple  crystals  of  the 
alum  are  deposited.  This  change  can  be  facilitated  by 
putting  some  crystals  of  the  salt  in  the  concentrated 
green  solution.  The  action  of  reducing  agents  upon  po- 
tassium dichromate  will  be  treated  of  farther  on.  The 
salt  finds  application  in  dyeing  and  tanning. 

Chromic  Acid  and  the  Chromates. — It  has  already  been 
stated  that  when  chromium  compounds  belonging  to  the 
chromous  and  chromic  series  are  oxidized  in  the  pres- 
ence of  bases  they  are  converted  into  chromates.  These 
salts  are  derived  from  an  acid  of  the  formula  H2CrO4, 
which  is  unknown,  as  it  breaks  down  spontaneously  into 
chromium  trioxide,  CrO3,  and  water,  when  it  is  set  free 
from  its  salts,  just  as  carbonic  and  sulphurous  acids  break 
down  respectively  into  carbon  dioxide  and  water,  and 
sulphur  dioxide  and  water.  The  starting-point  for  the 
preparation  of  the  chromates  and  the  compounds  re- 
lated to  them  is  chromic  iron.  This  is  ground  fine,  inti- 
mately mixed  with  a  mixture  of  caustic  potash  and  lime, 
and  then  heated  in  shallow  furnaces  in  contact  with  the 
air.  '  Under  these  circumstances  oxidation  is  effected  by 
the  oxygen  of  the  air.  The  iron  is  converted  into  ferric 
oxide,  and  the  chromium  gives,  with  the  calcium  and 
potassium,  the  corresponding  chromates,  CaCrO4  and 
K2CrO4.  When  the  mass  is  treated  with  water  these 
salts  dissolve,  and  ferric  oxide  remains  undissolved.  By 
treating  the  solution  with  potassium  sulphate  the  cal- 
cium salt  is  converted  into  the  potassium  salt,  and  thus 


664  INORGANIC  CHEMISTRY. 

all  the  chromium  appears  in  the  form  of  potassium  ehro* 
mate,  the  changes  referred  to  are  represented  in  the  fol- 
lowing equations : 

2(Cr02)2Fe  +  8KOH  +  7O  =  4K2CrO4  +  Fe2O3  +  4H2O  : 
2(CrO2)2Fe  +  4CaO    +  7O  =  4CaCrO4  +  Fe2O3 ; 
OaCr04       +  K2SO4  =  K2CrO4    +  CaSO4. 

As  potassium  chromate  is  easily  soluble  in  water,  and 
therefore  difficult  to  purify,  it  is  converted  into  the 
dichromate,  which  is  less  soluble  and  crystallizes  well. 
The  change  is  easily  effected  by  adding  the  necessary 
quantity  of  a  dilute  acid.  If  nitric  acid  is  used  the  re- 
action is  represented  by  the  following  equation : 

2K2Cr04  +  2HN03  =  K2Cr2O7  +  2KNO3  +  H2O. 

The  salt  thus  obtained  is  manufactured  on  the  large 
scale  and  is  the  starting-point  for  the  preparation  of 
other  chromium  compounds. 

Potassium  Chromate,  K2CrO4,  formed  as  above  de- 
scribed, is  a  light-yellow  crystallized  substance  which  is 
easily  soluble  in  water.  It  is  isomorphous  with  potas- 
sium sulphate.  Acids  convert  it  into  the  dichromate,  as 
just  stated. 

Potassium  Dichromate,  K2Cr2O7. — This  salt  forms  large 
red  plates,  which  are  triclinic.  It  is  soluble  in  ten  parts 
of  water  at  the  ordinary  temperature,  and  is  much  more 
soluble  in  hot  water.  When  heated,  it  at  first  melts 
without  undergoing  decomposition ;  at  white  heat,  how- 
ever, it  is  decomposed,  yielding  the  chromate,  chromic 
oxide,  and  oxygen  : 

2K2Cr2O7  =  2K2Cr04  +  Cr2O3  +  3O. 

It  undergoes  a  similar  change,  but  much  more  readily, 
when  heated  with  concentrated  sulphuric  acid.  In  this 
case,  however,  the  chromic  oxide  forms  chromic  sulphate 
with  the  acid,  and  this  forms  chrome-alum  with  the  po- 
tassium sulphate : 

K2Cr2O7  +  4H2SO4  =  2KCr(SO4)2  +  4H2O  +  3O. 
All  the  oxygen  in  the  chromate  in  excess  of  that  required 


POTASSIUM  DICHROMATE.  665 

to  form  the  alum  and  water  is  given  off.  This  also  is 
the  character  of  the  action  towards  reducing  agents  in 
general.  One  molecule  of  the  dichromate  gives  three 
atoms  of  oxygen.  With  sulphur  dioxide  the  action  is 
that  represented  in  the  equation 

K3O2O7  -f  4H2SO4  +  3SO2  =  2KCr(SO4)2  +  3H2SO4  -f  H2O. 

Or,  one  molecule  of  the  dichromate  converts  three  mole- 
cules of  sulphur  dioxide,  SO2,  into  three  molecules  of 
sulphuric  acid,  H2SO4. 

The  action  with  alcohol  will  be  understood  by  the  aid 
of  the  following  equation,  which  represents  the  action  of 
oxygen  in  general  upon  alcohol : 

C2H60  +  O  =  C2H40  +  H,0. 

Alcohol  Aldehyde 

Each  molecule  of  alcohol  requires  one  atom  of  oxygen 
to  convert  it  into  aldehyde.  Therefore,  one  molecule  of 
the  dichromate  oxidizes  three  molecules  of  alcohol  to 
aldehyde : 

K2Cr207  +  4H2SO4  +  3C2H6O  =  2KCr(SO4)2  +  3C2H4O  +  7H2O. 

Concentrated  hydrochloric  acid  is  oxidized  by  the  di- 
chromate, and  chlorine  is  evolved : 

K2Cr207  +  14HC1  =  2KC1  +  2CrCl3  +  7H2O  +  601. 

Here  two  atoms  of  chlorine  are  required  to  form  potas- 
sium chloride  with  the  potassium,  and  six  to  form  chro- 
mic chloride  with  the  chromium ;  and  the  eight  hydro- 
gen atoms  in  combination  with  this  chlorine  combine 
with  four  atoms  of  oxygen  of  the  dichromate,  leaving 
three  more  to  oxidize  hydrochloric  acid.  Consequently 
one  molecule  of  the  dichromate  sets  free  six  atoms  of 
chlorine  : 

3O  +  6HC1  =  3H2O  +  6C1. 

When  the  dichromate  in  solution  is  treated  with  po- 
tassium hydroxide,  its  color  changes  to  yellow,  in  con- 
sequence of  the  formation  of  the  chromate,  the  action 
taking  place  as  represented  in  this  equation : 

K2Cr2O7  +  2KOH  =  2K2CrO4  -f  H2O. 


6 GO  INORGANIC  CHEMISTRY. 

Potassium  dichromate  finds  extensive  use  in  the  arts 
and  in  the  laboratory  as  an  oxidizing  agent.  With  gela- 
tine it  forms  a  mixture  which  is  sensitive  to  light,  which 
turns  it  dark,  and  makes  it  insoluble.  This  fact  is  made 
the  basis  of  a  number  of  photographic  processes.  The 
dichromate  is  used,  further,  in  dyeing. 

Chromium  Trioxide,  CrO3,  crystallizes  out  on  cooling 
when  either  the  chromate  or  the  dichromate  is  treated 
in  concentrated  solution  with  concentrated  sulphuric  acid. 
This  is  a  beautiful  red  substance,  which  crystallizes  in 
needles.  When  dissolved  in  water  it  forms  a  solution 
from  which,  by  neutralization,  the  chromates  can  be  ob- 
tained. When  heated  alone  it  gives  off  half  its  oxygen, 
and  is  converted  into  chromic  oxide  : 

2Cr03  =  Cr203  +  3O  ; 

and  when  heated  with  sulphuric  acid  it  gives  chromic 
sulphate  and  oxygen  : 

2CrO3  +  3H2SO4  =  Cr2(SO4)3  +  3H2O  +  3O. 

It  is  an  extremely  active  oxidizing  agent,  disintegrating 
most  organic  substances  with  which  it  is  brought  in  con- 
tact. 

Relations  between  the  Chromates  and  Bichromates. — 
The  fact  that  chromium  trioxide  with  water  gives  chro- 
mic acid,  which  is  a  dibasic  acid,  whose  salts  in  general 
resemble  those  of  sulphuric  acid,  leads  to  the  belief  that 
the  structure  of  chromic  acid  should  be  represented  by 
a  formula  similar  to  that  of  sulphuric  acid,  thus  : 

O  O 

HO-S-OH  HO-Cr-OH 

6  b. 

or  or 

0,S(OH),  02Cr(OH), 

Just  as  sulphuric  acid  by  loss  of  water  is  converted  into 
disulphuric  acid  or  pyrosulphuric  acid,  so  chromic  acid 
is  converted  into  dichromic  acid,  and  in  all  probability 
the  relation  between  the  chromates  and  dichromates  is 


RELATIONS  OF  THE  CHROMATES  AND  BICHROMATES.     667 

the  same  as  that  between  the  sulphates  and  disulphates, 
as  represented  by  the  equations 

9Ot< 


CrO,< 

But,  as  has  been  stated,  neither  chromic  acid  nor  di- 
chromic acid  is  known,  as  they  break  down  into  chromium 
trioxide  and  water  when  set  free  from  their  salts.  The 
conversion  of  potassium  chromate  into  the  dichromate 
by  treatment  with  an  acid  is  represented  as  follows  : 


.OK  Prft    . 

'<OK      HNO,  _  Cr°><OH 

N°S  "  OrO.<g|+ 


CrO,<K 


Sodium  Chromate,  Na2CrO4,  and  Sodium  Dichromate, 
Na2Cr2O7,  are  made  in  the  same  way  as  the  potassium 
compounds.  They  are  both  deliquescent  as  ordinarily 
made.  It  has,  however,  recently  been  shown  that  a  so- 
dium dichromate  which  is  not  deliquescent  can  be  made  ; 
and  at  present  it  is  largely  manufactured  instead  of  the 
somewhat  more  expensive  potassium  salt. 

Barium  Chromate,  BaCrO4,  like  the  sulphate,  is  insol- 
uble in  water,  and  is  precipitated  when  a  solution  of  a 
barium  salt  is  brought  together  with  a  soluble  chromate 
or  dichromate.  It  is  soluble  in  hydrochloric  acid  and 
nitric  acid,  but  not  in  acetic  acid.  The  strontium  salt  is 
somewhat  soluble  in  water,  and  easily  in  hydrochloric,. 
nitric,  and  acetic  acids. 


668  INORGANIC  CHEMISTRY. 

Lead  Chromate,  PbCrO4,  occurs  in  nature,  and  is  formed 
as  a  beautiful  yellow  precipitate  when  a  solution  of  a 
lead  salt  is  treated  with  a  solution  of  a  chromate  or  di- 
chromate.  It  is  used  as  a  pigment  under  the  name 
chrome  yellow.  When  heated  to  a  high  temperature  it 
gives  off  some  of  its  oxygen.  In  contact  with  oxidizable 
substances  it  gives  off  its  oxygen  very  easily,  and  it  is 
used  in  the  laboratory  in  some  cases  instead  of  cupric 
oxide  in  the  analysis  of  organic  compounds  (see  p.  594). 
Treated  with  dilute  potassium  or  sodium  hydroxide  in 
insufficient  quantity  to  dissolve  it,  it  turns  red  in  conse- 
quence of  the  formation  of  a  basic  chromate  which  is 
known  as  chrome  red.  The  action  takes  place  as  repre- 
sented in  the  equation 

CrO,<°>Pb 

2KOH  = 


This  then  loses  water,  and  forms  the  salt  CrO2<Q~pi  >O, 

which  is  chrome  red. 

Lead  chromate  dissolves  completely  in  the  caustic 
alkalies  in  consequence  of  the  formation  of  chromates 
and  plumbites. 

Silver  chromate,  Ag2CrO4,  is  formed  as  a  red  precipitate 
when  a  chromate  is  treated  with  a  silver  salt. 

Potassium  trichromatet  K2Cr3O]0,  and  potassium  tetra- 
cJiromate,  K2O4O13,  are  formed  by  treating  the  dichromate 
with  nitric  acid  : 

3K2Cr2O7  +  2HNO3  =  2K2Cr3O10  +  2KN03  +  H2O  ; 
2K2Cr20,  +  2HN03  -  K2Cr4O13    +  2KNO3  +  H2O. 

The  acids  from  which  these  salts  are  derived  bear  to  or- 
dinary chromic  acid  relations  similar  to  those  which  the 
polysilicic  acids  bear  to  ordinary  silicic  acid. 

Chromium  Oxychloride,  Chromyl  Chloride,  CrOzCla, 
is  analogous  to  sulphuryl  chloride,  SO2C12,  and  is  to  be 


ANALYTICAL  REACTIONS  OF  CHROMIUM.  669 

regarded  as  derived  from  chromic  acid  by  the  replace- 
ment of  the  hydroxyls  by  chlorine  : 

CrO.<g|;  OrO.<g|.    /  -; 

It  is  formed  by  treating  a  mixture  of  sodium  chloride 
and  potassium  dichromate  with  sulphuric  acid.  Prob- 
ably the  sulphuric  acid  sets  free  hydrochloric  acid  and 
chromic  acid  simultaneously,  and  these  then  act  upon 
each  other  as  represented  in  the  equation 


It  is  a  dark  -red-colored  liquid,  which  boils  at  116°  with- 
out decomposition.  With  water  it  gives  chromic  acid 
and  hydrochloric  acid,  and  with  an  alkaline  hydroxide  it 
gives  the  corresponding  chromate  : 


182  =  Cr°°<  2HCL 


Reactions  which  are  of  Special  Value  in  Chemical  Anal- 
ysis. —  The  reactions  of  chromic  salts  with  the  alkaline 
hydroxides  have  been  explained.  With  the  soluble  car- 
bonates they  give  precipitates  which  consist  mainly  of 
the  hydroxide,  though  carbonic  acid  is  to  some  extent 
in  combination  in  them.  These  precipitates  are  soluble 
in  a  large  excess  of  the  carbonate. 

The  solution  of  chromic  oxide  in  an  alkali  is  green, 
but  by  oxidizing  agents  it  is  turned  yellow  in  consequence 
of  the  formation  of  a  chromate. 

Hydrogen  sulphide  does  not  precipitate  chromium  from 
its  salts.  Ammonium  sulphide  precipitates  chromium  hy- 
droxide, the  reaction  being  the  same  as  in  the  case  of 
aluminium  (which  see). 

Chromates  give  with  barium  and  lead  salts  yellow  pre- 
cipitates (see  Barium  Chromate  and  Lead  Chromate). 


670  INORGANIC  CHEMISTRY. 

When  heated  with  sulphuric  acid,  the  chromates  are 
decomposed,  with  evolution  of  oxygen.  The  action  of 
hydrochloric  acid  upon  the  chromates  was  explained 
above. 

In  neutral  or  slightly  acid  solutions  of  chromates, 
hydrogen  sulphide  and  ammonium  sulphide  act  as  re- 
ducing agents,  and  precipitate  the  hydroxide  mixed  with 
sulphur : 

K2Cr2O7+  3HaS+2HCl  =  2KC1  +  2Cr(OH)3+3S+  H2O. 

With  borax  or  microcosmic  salt  chromium  compounds 
form  green  beads,  both  in  the  oxidizing  and  reducing 
flames. 

MOLYBDENUM,  Mo  (At.  Wt.  95.26). 

General. — Molybdenum  is  of  interest  on  account  of  the 
variety  of  its  compounds.  It  forms  four  compounds  with 
chlorine,  the  formulas  of  which  appear  to  be  MoCla, 
MoCl3,  MoCl4,  and  MoCl5.  With  oxygen  also  it  forms 
four  compounds,  but,  while  the  first  three  are  analogous 
to  the  first  three  chlorine  compounds  above  mentioned, 
the  last  differs  from  the  last  chlorine  compound.  In  it 
the  element  is  sexivalent.  The  formulas  of  the  oxygen 
compounds  are  MoO,  Mo2O3,  MoO2,  and  MoO3.  The 
last  of  these  is  analogous  to  chromium  trioxide,  as  it 
forms  salts  with  bases  analogous  to  the  chromates,  and 
known  as  the  molybdates. 

Occurrence  and  Preparation. — Molybdenum  occurs  in 
nature  principally  as  molybdenite,  which  is  the  sulphide 
MoS2,  and  as  wulfenite,  which  is  lead  molybdate,  PbMoO4. 
It  occurs  also,  but  in  smaller  quantity,  as  molybdenum 
trioxide,  MoO3.  It  is  obtained  in  free  condition  by  heat- 
ing the  oxides  or  chlorides  in  a  current  of  hydrogen. 

Properties. — It  is  a  very  hard  silver-white  metal  of 
specific  gravity  8.6.  It  is  apparently  infusible  when  pure. 
When  heated  for  a  considerable  time  in  the  air  it  is  con- 
verted into  the  trioxide.  It  dissolves  in  concentrated 
nitric  acid  and  in  aqua  regia. 

Chlorides. — When  molybdenum  is  heated  for  a  long 
time  in  a  current  of  chlorine  it  is  finally  completely  con- 
verted into  the  pentachloride,  MoCl5.  When  the  penta- 


OXIDES  OF  MOLYBDENUM.  671 

chloride  is  heated  to  250°  in  a  current  of  hydrogen  it  is 
converted  into  the  trichloride,  MoCl3.  And  when  the 
trichloride  is  heated  in  a  current  of  carbon  dioxide  it  is 
converted  into  the  dichloride  and  the  tetrachloride : 

2MoCl3  =  MoCla  +  MoCl4. 

Oxides. — The  final  product  of  the  action  of  oxygen  or 
oxidizing  agents  upon  molybdenum  is  the  trioxide,  MoO3. 
This  is  obtained  comparatively  easily  when  molybdenite, 
MoS2,  is  mixed  with  pure  sand  and  roasted,  then  treated 
with  ammonia,  which  forms  ammonium  molybdate.  This 
salt  is  crystallized,  and  afterwards  decomposed  by  nitric 
acid.  The  oxide  is  a  white  crystallized  substance  which 
is  difficultly  soluble  in  water.  When  a  solution  of  the 
trioxide  is  treated  with  reducing  agents,  as  sodium-amal- 
gam, lower  oxides  are  formed,  and  the  reduction  stops 
at  molybdic  oxide,  Mo2O3.  This  is  a  black  compound. 
The  dioxide,  MoO2,  is  a  dark-blue  substance  which  is 
formed  by  gently  heating  the  metal  or  molybdic  oxide, 
Mo2O3,  in  the  air,  and  also  by  heating  the  corresponding 
hydroxide,  Mo(OH)4. 

Molybdenum  monoxide,  MoO,  is  formed  by  treating  a 
solution  of  the  dichloride  with  hot  caustic  potash.  Like 
molybdic  oxide  it  is  black. 

Molybdenum  trisulphide,  MoS3,  is  precipitated  as  a  red- 
dish-brown substance  when  a  moderately  concentrated 
solution  of  a  molybdate  is  treated  with  hydrogen  sul- 
phide. When  heated  it  is  converted  into  the  disulphide, 
MoS2,  which  is  identical  in  composition  with  the  molyb- 
denite found  in  nature. 

Molybdic  Acid  and  the  Molybdates. — Molybdenum  tri- 
oxide, MoO3,  combines  readily  with  bases,  forming  salts 
which  in  composition  are  analogous  to  the  chromates. 
When  ammonium  molybdate  is  treated  with  moderately 
dilute  nitric  acid  free  molybdic  acid  crystallizes  out. 
This  has  the  composition  represented  by  the  formula 

roH 

f  OTT 
H2MoO4  +  H2O,  and  perhaps  the  constitution  MoO  -j  QTT, 

[OH 


672  INORGANIC   CHEMISTRY. 

analogous  to  that  of  tetrahydroxyl-sulphuric  acid.  Mo- 
lybdic acid  not  only  forms  compounds  with  bases,  but 
also  with  acids.  The  latter  are,  however,  not  salts,  but 
complex  acids  which  in  turn  form  salts  with  bases,  some 
of  which  are  extremely  complicated.  There  are  a  num- 
ber of  molybdates  known  of  the  general  formula  M2MoO4, 
in  which  M  represents  a  univalent  element,  but  the  acid 
further  shows  a  marked  tendency  to  form  more  complex 
salts,  derivatives  of  polymolybdic  acids.  Thus  the  fol- 
lowing sodium  salts  are  known  :  Na2MoO4,  Na2Mo2O7, 
Na2Mo3010,  Na2Mo4013,  Na2Mo8O25,  Na2Mo10O31,  and 
Na,Mo7O24.  The  relations  between  the  acids  from  which 
these  salts  are  derived,  and  molybdic  acid,  H2MoO4,  will 
readily  be  seen  by  the  aid  of  the  following  equations : 

2HJMo04  =  H2Mo2O7  +  H2O  ; 
3H2Mo04  =  H2Mo3010  +  2H2O ; 
4H2Mo04  =  H2Mc4013  +  3H2O ; 
8H3MoO4  =  H2Mo802B  +  7H2O ; 
10H2Mo04  =  H2Mo10031  +  9H20  ; 
7H2Mo04  =  H6Mo7O2,  +  4H2O. 

Lead  Molybdate,  PbMoO4,  occurs  in  nature,  as  has  been 
stated,  and  the  mineral  is  known  as  wulfenite.  It  can  be 
obtained  artificially  by  melting  together  sodium  molyb- 
date,  lead  chloride,  and  sodium  chloride  ;  or  by  treating 
a  solution  of  sodium  molybdate  with  a  solution  of  lead 
nitrate.  If  the  reagents  are  pure  the  artificially  prepared 
salt  is  white,  while  the  natural  variety  is  always  yellow 
or  red. 

Phospho-molybdic  Acid. — Among  the  best  known  and 
most  frequently  met  with  compounds  of  molybdic  acid 
with  acids  is  that  which  it  forms  with  phosphoric  acid, 
known  as  phospho-molybdic  acid.  When  a  solution  of 
ammonium  molybdate  in  an  excess  of  nitric  acid  is  added 
in  excess  to  a  solution  of  phosphoric  acid  or  a  phosphate, 
a  yellow  precipitate  is  formed.  This  is  ammonium  phos- 
pho-molybdate,  which,  when  dried,  has  the  composition 
represented  by  the  formula  12MoO3.(NH4)3PO4.  This  is 
insoluble  in  water  and  in  dilute  acids,  and  also  in  a  nitric- 


TUNGSTEN.  673 

acid  solution  of  ammonium  molybdate.  On  account  of 
the  properties  mentioned,  this  salt  furnishes  a  valuable 
means  of  detecting  phosphoric  acid  and  of  precipitating 
it  from  its  solutions.  When  the  salt  is  treated  with  aqua 
regia  it  is  decomposed,  and  from  the  solution  formed 
a  compound  of  the  composition  H3PO4.llMoO3-|- 12H2O 
crystallizes  out. 

TUNGSTEN,  W  (At.  Wt.  183.43). 

General. — Like  molybdenum,  tungsten  forms  a  large 
variety  of  compounds.  With  chlorine  it  forms  four,  of 
which  the  formulas  are  WC12,  WC14,  WC16,  and  WC16. 
With  oxygen,  however,  it  forms  but  two  compounds,  and 
these  are  represented  by  the  formulas  WO2  and  WO3. 
The  trioxide  forms  salts  with  bases  which  are  analogous 
to  the  molybdates,  and,  like  molybdic  acid,  tungstic  acid 
forms  complicated  salts  which  are  derived  from  poly- 
tungstic  acids.  Further,  tungstic  acid  combines  with 
other  acids,  forming  very  complex  acids. 

Occurrence  and  Preparation. — Tungsten  occurs  in  na- 
ture as  tungstates.  The  principal  one  is  the  iron  salt, 
which  always,  however,  contains  some  manganese.  This 
is  known  as  wolframite,  and  has  the  composition  repre- 
sented by  the  formula  FeWO4.  Calcium  tungstate, 
CaWO4,  or  scheelite,  and  lead  tungstate,  PbWO4,  or 
stolzite,  are  also  found  in  nature,  but  in  smaller  quantity 
than  wolframite.  The  element  is  prepared  by  reducing 
the  chlorides  or  oxides  in  a  current  of  hydrogen. 

Properties. — Tungsten  forms  lustrous,  steel-colored 
laminae,  or  a  black  powder.  It  is  very  hard  and  difficultly 
fusible,  and  has  the  specific  gravity  19.129.  It  is  not 
changed  by  contact  with  the  air  at  ordinary  temperatures. 
At  higher  temperatures  it  combines  with  oxygen  and 
forms  the  trioxide,  WO3.  Nitric  acid  and  aqua  regia 
convert  it  into  the  trioxide.  It  is  used  in  the  manufac- 
ture of  steel,  as  the  addition  of  from  8  to  9  per  cent  of  it 
makes  steel  extremely  hard. 

Chlorides. — When  tungsten  is  heated  in  a  current  of 
chlorine  it  is  converted  into  the  hexachloride,  WC16.  The 
other  chlorides  are  formed  by  heating  this  in  hydrogen. 


674  INORGANIC  CHEMISTRY. 

Oxides. —  Tungsten  trioxide,  WO3,  is  found  in  small 
quantity  in  nature,  and  is  formed  from  wolframite  by  a 
number  of  methods.  When  a  solution  of  a  tungstate  is 
boiled  with  an  acid  the  trioxide  is  precipitated  as  a  yel- 
low powder.  Under  the  influence  of  sunlight  it  turns 
greenish.  It  is  insoluble  in  water.  In  alkalies  it  dis- 
solves, forming  the  tungstates.  When  heated  in  a  current 
of  hydrogen,  lower  oxides  are  formed ;  and  a  blue  com- 
pound thus  obtained  appears  to  have  the  composition 
represented  by  the  formula  2WoO3  -[-  WO2  or  W3O8. 
The  dioxide,  WO2,  is  obtained  by  further  reduction. 
This  is  a  brown  powder. 

Tungstic  Acid  and  the  Tungstates. — When  the  required 
quantities  of  tungsten  trioxide  and  potassium  carbonate 
are  brought  together  in  solution,  or  are  melted  together, 
potassium  tungstate,  K2WO4,  is  formed.  If  a  solution  of 
this  salt  is  treated  with  a  strong  acid  at  the  ordinary 
temperature,  a  white  precipitate  of  the  composition 
H2WO4-f-H2O  is  formed.  This  is  tungstic  acid,  analogous 
to  crystallized  molybdic  acid.  If  the  solutions  are  hot 
the  precipitate  has  the  composition  H2WO4.  Among  the 
complex  salts  derived  from  polytungstic  acids  are  the 
following  :  Na2W2O7,  Na4W3On,  Na10W]2O41,  etc.  The  re- 
lations between  the  polytungstic  acids  and  the  ordinary 
variety  of  the  acid,  H2WO4,  will  be  readily  understood. 
The  salt,  Na10W12O41  -j-  28H2O,  is  known  as  sodium  para- 
tungstate.  It  is  manufactured  on  the  large  scale  by  heat- 
ing together  wolframite  and  calcined  sodium  carbonate. 
Inflammable  substances  impregnated  with  a  solution  of 
the  salt  burn  with  great  difficulty,  and  it  is  used  to  pro- 
tect various  articles  from  fire.  The  salts  derived  from 
the  acid,  H2W4O13,  are  called  metatung 'states. 

Silico-tungstic  Acids. — Among  the  most  interesting  of 
the  complex  compounds  formed  by  the  combination 
of  tungstic  acid  with  other  acids  are  those  known  as 
the  silico-tungstic  acids.  When  sodium  paratungstate, 
Na10W12O41,  is  boiled  in  solution  with  precipitated  gelati- 
nous silicic  acid,  the  latter  dissolves,  and  from  the  solution 
a  salt  of  the  composition  Na8W12SiO42  +  7H2O  crystal- 
lizes. This  is  soluble  in  one  fifth  its  weight  of  water, 


URANIUM.  675 

and  the  solution  has  the  remarkably  high  specific  gravity 
3.05.  The  acid  from  which  the  salt  is  derived  is  known 
as  silico-tungstic  acid.  Its  composition  is  represented 
by  the  formula  4H2O.12WO3.SiO2 ;  and  it  may  be  re- 
garded as  made  up  of  a  polytungstic  acid,  H^W^O^, 
in  combination  with  one  molecule  of  silicon  dioxide, 
H8W1204,SiO, 

UBANIUM,  U  (At.  Wt.  237.77). 

General. — Uranium  has  stronger  basic  properties  than 
either  molybdenum  or  tungsten ;  and  it  differs  from 
chromium  in  the  fact  that  the  trioxide  forms  salts  with 
acids.  These  salts  are  the  uranyl  salts  which  are  derived 
from  the  hydroxide,  UO2(OH)2  or  H2UO4.  The  cor- 
responding compounds  of  chromium,  molybdenum,  and 
tungsten  are  acids.  Uranium  also  forms  salts  in  which 
it  acts  as  a  quadrivalent  element,  as  U(SO4)2.  While  the 
hydroxide,  UO2(OH)2,  forms  salts  with  acids,  it  also  forms 
salts  with  the  strongest  bases.  These  are  analogous  in 
composition  to  the  dichromates,  and  have  the  general 
formula  M2U2O7.  With  chlorine,  uranium  forms  the 
compounds  UC13,  UC14,  and  UC15 ;  and  with  oxygen  the 
following :  U02,  U3O8,  UO3,  and  UO4. 

Occurrence  and  Preparation. — Uranium  occurs  in  na- 
ture chiefly  in  the  form  of  the  mineral  known  as  pitch- 
blende or  uraninite,  which  consists  of  the  oxide,U3O8,  mixed 
with  a  number  of  other  substances  in  smaller  or  larger 
quantities.  When  this  is  finely  powdered  and  treated 
with  concentrated  nitric  acid,  uranyl  nitrate,  UO2(NO3)2, 
is  obtained,  and,  by  igniting  this,  the  trioxide,  UO3,  is 
left  behind.  In  order  to  isolate  the  metal,  the  oxide  thus 
obtained  is  mixed  with  charcoal  and  treated  with  chlo- 
rine, when  the  tetrachloride,  UC14,  is  formed.  This  is 
then  reduced  by  heating  it  with  sodium  under  a  cover  of 
the  molten  chlorides  of  potassium  and  sodium. 

Properties. — Uranium  has  the  color  of  nickel  and  the 
specific  gravity  18.4.  When  heated  to  redness  it  is  oxid- 
ized superficially.  It  dissolves  in  dilute  acids  with 
evolution  of  hydrogen. 


676  INORGANIC  CHEMISTRY. 

Chlorides.  —  When  chlorine  acts  upon  finely  divided 
uranium,  the  two  combine  to  form  the  tetrachloride,  UC14. 
When  this  is  heated  in  hydrogen  it  loses  a  part  of  its 
chlorine  and  forms  the  trichloride,  UC13  ;  and  when  the 
tetrachloride  is  treated  with  chlorine  it  is  partly  converted 
into  the  pentachloride,  UC15.  The  tetrachloride  is  the 
most  stable  form. 

Oxides.  —  The  oxide  of  uranium  which  is  formed  as  the 
last  product  of  the  action  of  oxygen  on  uranium  or  the 
other  oxides  when  these  are  heated  in  the  air  is  that  which 
has  the  composition  U3O8,  which  is  also  the  composition 
of  the  natural  variety.  When  this  is  treated  with  nitric 
acid,  however,  it.  is  converted  into  uranyl  nitrate, 
TJO2(NO3)2,  which  is  a  derivative  of  the  trioxide,  UO3  ; 
and  when  the  nitrate  is  ignited,  the  trioxide  is  left  behind. 
By  reduction  with  hydrogen  the  trioxide  is  converted  into 
the  dioxide,  UO2  ;  and  when  either  the  dioxide  or  the  tri- 
oxide is  heated  in  the  air,  the  product  obtained  is  the 
oxide  U3O8.  As  will  be  pointed  out  below,  this  is  re- 
garded as  a  uranium  salt  of  uranic  acid. 

Uranous  Salts.  —  In  the  uranous  salts,  uranium  acts  as 
a  quadrivalent  element,  replacing  four  atoms  of  hydro- 
gen, as,  for  example,  in  the  sulphate,  which  has  the  com- 
position U(SO4)2.  But  few  salts  of  this  order  are  known. 

Uranyl  Salts.  —  As  already  explained,  the  uranyl  salts 

OH 

are    derivatives   of   the   hydroxide   UO2<QTT,  and   are 

formed  by  the   action  of   acids,  as  represented  in  the 
equations  below  : 

uo.<81  +  HO>S0'  =  UO,<g>SO,  +  2H,0; 


TTO   ,         ,   HO.NO,       TTO    .O.NO, 
U05<OH 


They  are  derived  from  the  acids  by  replacing  the  hydro- 
gen by  uranyl,  UOa,  which  is  bivalent.  —  Uranyl  nitrate, 
TJO2(NO3)3,  is  easily  obtained,  as  above  described,  and 
crystallizes  well  in  lemon-yellow  prisms.  —  Uranyl  sul- 
phate, UO3(SO4),  is  formed  by  treating  the  nitrate  with 


URANATES.  677 

sulphuric  acid.  It  combines  with  ammonium  sulphate, 
forming  the  salt  UO2(SO4)  +  (NH4)2SO4  +  2H2O,  which 
is  difficultly  soluble  in  water  and  crystallizes  in  lemon- 
yellow  prisms. 

Uranates. — When  a  uranyl  salt  is  treated  with  a  soluble 
hydroxide  a  precipitate  is  formed  which  is  a  salt  of  an 
acid,  H2U2O7,  which  may  be  called  diuranic  acid,  as  in 
composition  it  is  analogous  to  dichromic  and  disulphuric 
acids. 

Sodium  diuranate,  Na2U2O7,  is  a  fine  yellow  powder,  and 
is  manufactured  and  sold  under  the  name  uranium  yellow, 
being  used  as  a  pigment  for  coloring  glass,  etc. — Ammo- 
nium diuranate,  (NH4)2U2O7,  is  also  manufactured  on  the 
large  scale.  When  it  is  treated  with  a  solution  of  am- 
monium carbonate  it  dissolves,  and  from  the  solution  a 
salt  of  the  composition  UO2(CO3)  +  2(NH4)2CO3  crystal- 
lizes out.  The  solubility  of  ammonium  diuranate  in  am- 
monium carbonate  is  utilized  in  analysis. — The  oxide  of 
the  formula  U3O8  above  referred  to  may  be  a  uranous 
salt  of  uranic  acid  as  represented  by  tbe  formula 


0 1 

UO,<r\ 

§k 
uo,<<>j 


Many  uranium  salts  exhibit  in  solution  a  beautiful 
fluorescence. 


CHAPTER  XXXII. 

ELEMENTS  OF  FAMILY  VII,   GROUP  A : 
MANGANESE  (Mn,  At.  Wt.  54.57). 

General. — At  the  close  of  Chapter  XII  (which  see), 
which  treated  of  the  elements  of  Family  VII,  Group 
B,  or  the  chlorine  group,  reference  was  made  to  man- 
ganese, and  attention  was  called  to  the  fact  that  in  some 
respects  it  resembles  chlorine.  The  resemblance  is 
seen  in  the  formation  of  an  oxide,  Mn2O7,  and  an  acid, 
HMnO4,  analogous  to  perchloric  acid.  On  the  other 
hand,  in  many  of  its  compounds  it  plays  the  part  of  a 
base-forming  element,  and  in  this  capacity  it  forms  two 
series  of  compounds,  known  as  the  manganous  and  the 
manganic  compounds.  In  the  former  the  element  ap- 
pears to  be  bivalent,  and  in  the  latter  trivalent.  The 
formulas  of  some  of  the  principal  manganous  com- 
pounds are  : 

MnCl2,  Mn(OH)2,  MnO,  Mn(N08)2,  MnSO4,  MnCO3,  etc. 

The  formulas  of  some  of  the  principal  manganic  com- 
pounds are : 

Mn(OH)3,  MnaO3,  MnO(OH),  Mn2(SO4)s,  KMn(SO4)3  +  12H2O,  etc. 

These  two  series  of  compounds  are  analogous  in  composi- 
tion to  the  chromous  and  chromic  compounds,  but,  while 
the  chromic  compounds  are  more  stable  than  the  chro- 
mous compounds,  the  manganous  compounds  are  more 
stable  than  the  manganic  compounds.  By  contact  with 
the  air,  the  manganous  are  not  as  a  rule  converted  into 
the  manganic  compounds. 

Corresponding  to  chromic  acid,  there  is  a  manganic 
acid,  H2MnO4 ;  and,  further,  there  is  the  permanganic 
acid  already  mentioned,  of  the  formula  HMnO4.  An  anal- 
ogous compound  of  chromium,  perchromic  acid,  HCrO4, 

(678) 


MANGANESE.  679 

is  believed  to  be  formed  when  hydrogen  peroxide  is 
added  to  an  aqueous  solution  of  chromic  acid,  but  this  is 
by  no  means  certain.  Manganic  acid  and  its  salts  are 
very  unstable,  and  are  readily  converted  into  perman- 
ganic acid  and  the  permanganates.  On  the  other  hand, 
perchromic  acid,  if  it  exists  at  all,  is  spontaneously  de- 
composed, yielding  ordinary  chromic  acid.  To  sum  up, 
then,  both  chromium  and  manganese  form  four  classes 
of  compounds.  But,  while  chromium  under  ordinary 
circumstances  preferably  forms  chromic  salts  and  salts 
of  chromic  acid,  manganese  preferably  forms  manganous 
salts  and  salts  of  permanganic  acid. 

Manganese  forms  a  number  of  oxides  corresponding 
to  the  formulas  MnO,  Mn2O3,  Mn3O4,  MnO2,  and  Mn3O7. 
Of  these  probably  the  one  of  the  composition  Mn3O4 
and  perhaps  that  of  the  composition  Mn2O3  are  com- 
pounds of  the  others,  as  will  be  pointed  out  further  on. 

Forms  in  which  Manganese  occurs  in  Nature. — The 
principal  natural  compound  of  manganese  is  the  black 
oxide  or  pyrolusite,  MnO2.  Besides  this,  however,  there 
are  several  compounds  found  in  nature,  the  principal 
ones  being,  braunite,  Mn2O3,  hausmannite,  Mn3O4,  man- 
ganite,  Mn2O2(OH)2,  and  rhodocroisite,  which  is  the  car- 
bonate, MnCO3. 

Preparation  and  Properties. — The  metal  is  isolated 
from  its  oxides  by  heating  them  to  a  high  temperature 
with  charcoal.  It  looks  like  cast  iron,  is  brittle  and 
hard,  and  has  the  specific  gravity  8.  It  easily  becomes 
oxidized  in  the  air,  decomposes  warm  water,  and  dis- 
solves readily  in  dilute  acids.  It  is  used  as  a  con- 
stituent of  some  useful  alloys,  and  imparts  certain  de- 
sirable properties  to  iron,  as  will  be  pointed  out  when 
that  metal  is  taken  up. 

Manganous  Chloride,  MnCl2. — This  chloride  is  obtained 
in  solution  by  dissolving  any  one  of  the  oxides  or  hy- 
droxides or  the  carbonate  of  manganese  in  hydrochloric 
acid  with  the  aid  of  gentle  heat.  Its  formation  in  the 
preparation  of  chlorine  from  manganese  dioxide  and  hy- 
drochloric acid  was  referred  to  under  Chlorine  (which 
see).  When  the  solution  is  evaporated  to  the  proper 


680  INORGANIC  CHEMISTRY. 

concentration,  the  salt  crystallizes  out  in  pink,  mono- 
clinic  plates  of  the  composition  MnCl2  -f-  4H2O.  When 
the  crystallized  salt  is  heated,  it  decomposes  into  the 
oxide  and  hydrochloric  acid,  as  so  many  other  chlorides 
do.  It  forms  double  chlorides,  an  example  of  which  is 
the  ammonium  compound  of  the  formula  MnCl2.2NH4Cl 
or  (NH4)2MnCl4. 

When  manganic  hydroxide,  Mn(OH)3,  is  treated  in  the 
cold  with  hydrochloric  acid,  a  deep  brown-colored  solu- 
tion is  formed,  which  is  believed  to  contain  the  trichlo- 
ride, MnCl3.  On  standing,  however,  this  solution  gives 
off  chlorine  slowly,  and  when  heated  it  gives  it  off  rap- 
idly, and  the  color  changes  to  pink  when  only  mangan- 
ous  chloride  is  left  in  the  solution. 

Phenomena  similar  to  those  just  mentioned  are  ob- 
served when  manganese  dioxide  is  treated  with  hydro- 
chloric acid  in  the  cold,  and  it  is  believed  that  the 
tetrachloride,  MnCl4,  is  contained  in  the  solution. 

While  the  tetrafluoride  itself  has  not  been  isolated,  a 
solution  is  obtained  by  treating  manganese  dioxide  with 
concentrated  hydrofluoric  acid  which  with  potassium 
fluoride  gives  a  salt  of  the  formula  MnF4.2KF  or 
K2MnF6. 

General  Remarks  Concerning  the  Oxides. — The  series  of 
oxides  of  manganese  strongly  suggests  that  of  the  oxides 
of  lead.  Manganese,  however,  forms  one  oxide,  the 
heptoxide,  Mn2O7,  for  which  there  is  no  analogue  among 
the  compounds  of  lead.  Placing  the  formulas  of  the 
oxides  of  the  two  metals  side  by  side,  we  have  the  follow- 
ing table : 

PbO  MnO 

Pb203  Mn203 

Pb3O4  Mn8O4 

PbO2  MnO2 

Mn207. 

Just  as  lead  oxide  when  heated  is  converted  into  red- 
lead,  Pb3O4,  so  the  other  oxides  of  manganese  are  con- 
verted into  the  oxide,  Mn3O4,  when  heated.  This  has  al- 
ready been  seen  in  the  preparation  of  oxygen  by  heating 
the  dioxide.  In  the  same  way,  the  oxide,  Mn2O3,  loses 


OXIDES  OF  MANGANESE.  681 

enough  oxygen,  and   the  lowest  oxide,  MnO,  takes   up 
enough  to  form  the  same  product : 

3Mn2O3  =  2Mn3O4  +  O  ; 
3MnO  +  O  =  Mn3O4. 

When  treated  with  energetic  oxidizing  agents  in  the 
presence  of  the  alkalies,  all  the  oxides  are  converted  into 
manganates.  On  the  other  hand,  if  a  manganate  is  re- 
duced in  the  presence  of  an  acid  the  tendency  to  form 
inanganous  compounds  shows  itself,  and  all  oxygen  pres- 
ent in  excess  of  that  required  to  form  the  manganous 
salt  is  given  off. 

Manganous  Oxide,  MnO,  is  formed  by  reducing  one  of 
the  higher  oxides  in  a  current  of  hydrogen. 

Manganous  Hydroxide,  Mn(OH)2,  is  formed  as  a  white 
precipitate  when  a  soluble  hydroxide  is  added  to  a  solu- 
tion of  a  manganous  salt.  Suspended  in  the  alkali,  or 
in  contact  with  air,  it  absorbs  oxygen,  and  is  converted 
into  hydroxides  corresponding  to  the  higher  oxides. 

Manganous-manganic  Oxide,  Mn3O4,  occurs  in  nature  as 
the  mineral  hausmannite.  It  is  formed,  as  already  stated, 
by  igniting  the  other  oxides  in  contact  with  the  air. 
When  heated  with  dilute  nitric  acid,  it  acts  like  the  cor- 
responding oxide  of  lead,  giving  manganous  nitrate,  and 
leaving  manganese  dioxide : 

Mn304  +  4HN03  =  2Mn(NO3)2  +  MnO,  +  2H2O. 

It  breaks  down  in  the  same  way  with  dilute  sulphuric 
acid.  These  facts  make  it  appear  probable  that  the  ox- 
ide is  the  manganous  salt  of  normal  manganous  acid, 
Mn(OH)4,  just  as  minium  or  red  lead  is  regarded  as  the 
lead  salt  of  normal  plumbic  acid,  Pb(OH)4,  (p.  651.)  This 

fg>Mn 

view  is  expressed  by  the  formula  Mn  i  JJ          .    The  de- 

[0>Mn 

composition  with  acids,  as  with  sulphuric  acid,  would, 
according  to  this,  be  represented  thus  : 

Mn04Mna  +  2H2SO4  =  2MnSO4  +  Mn(OH)4. 
The.  hydroxide  thus  formed  would  then  break  down  into 
the  dioxide  and  water. 


682  INORGANIC  CHEMISTRY. 

Manganic  Oxide,  Mn2O3,  occurs  in  nature  as  the  mm- 
eral  braunite,  and  it  can  be  made  from  the  other  oxides 
by  igniting  them  in  oxygen.  A  hydroxide  related  to 
this,  and  having  the  composition  MnO.OH,  analogous  to 
the  compounds  of  aluminium  and  chromium  of  the 
formulas  A1O.OH  and  CrO.OH,  is  found  in  nature,  and 
is  known  as  manganite.  The  hydroxide,  Mn(OH)3,  is 
formed  when  manganous  hydroxide,  Mn(OH)2,  is  ex- 
posed in  a  solution  of  ammonia  in  contact  with  the  air, 
and  forms  a  brownish  black  powder. 

Manganese  Dioxide,  MnO2.  —  This  important  compound 
occurs  in  nature  in  very  considerable  quantities,  and  is 
known  as  pyrolusite  or  the  black  oxide  of  manganese. 
It  is  obtained  artificially  by  gently  igniting  manganous 
nitrate.  A  hydroxide  derived  from  the  dioxide  is  ob- 
tained by  treating  a  manganous  salt  in  alkaline  solution 
with  a  soluble  hypochlorite  or  chlorine  or  bromine.  The 
chief  application  of  the  dioxide  is  in  the  preparation  of 
chlorine,  for  which  purpose  it  is  used  in  large  quantities. 
It  is  also  used  for  making  oxygen,  and  for  the  purpose  of 
decolorizing  glass.  In  the  last  process  a  small  quantity 
is  added  to  the  molten  glass.  This  alone  would  give  the 
glass  an  amethyst  color.  Without  it  the  glass  would 
be  green.  One  color  counteracts  the  other,  and  the  glass 
appears  colorless.  As  regards  the  action  of  hydro- 
chloric acid  upon  manganese  dioxide,  it  has  been  sug- 
gested, upon  the  basis  of  experimental  investigations, 
that  the  first  product  of  the  action  is  a  compound  of 
the  formula  H3MnCle,  which  is  the  chlorine  compound 
analogous  to  the  oxygen  acid,  HaMnO3.  The  action  is 
supposed  to  take  place  as  represented  in  the  following 
equation  : 

0=Mn=O       6HC1  =      ~>Mn=Cl       2HO. 


The  suggestion  is  made,  further,  that  it  is  this  compound, 
and  not  manganese  tetrachloride,  MnCl4,  which  breaks 
down  yielding  chlorine,  the  action  taking  place  thus  : 


H-(Cn>MnC1*  =  MnC1*  +  2HC1 


MANGANITES.  683 

The  manganons  chloride  and  some  of  the  chlormangan- 
ous  acid  then  react,  forming  a  compound  which  with 
water  undergoes  decomposition.  In  regard  to  this  sug- 
gestion, it  can  only  be  said  that  as  yet  it  is  not  sufficiently 
supported  by  facts.  The  formation  of  the  unstable 
compound,  H2MnCl6,  appears  highly  probable,  however, 
in  view  of  the  conduct  of  so  many  other  chlorides  in 
the  presence  of  hydrochloric  acid. 

Manganites. — There  are  some  salts  known  as  the  man- 
ganites, which  are  clearly  derived  from  hydroxides  re- 
lated to  manganese  dioxide.  Theoretically  the  simplest 
hydroxides  of  this  kind  are  those  of  the  formulas 
Mn(OH)4  and  MnO(OH)2.  The  salts  are  not,  however, 
derived  from  these,  but  from  more  complicated  forms,  as 
H2Mn2O6  and  H2Mn5Ou,  the  relations  between  which  and 
the  simpler  hydroxides  are  shown  in  the  equations 

2MnO(OH)2  =  Mn2O3(OH)2  +  H2O  ; 
5MnO(OH)2  =  Mn5O9(OH)2  +  4H2O. 

The  potassium  salt,  K2Mn5On,  is  obtained  when  carbon 
dioxide  is  conducted  into  a  solution  of  potassium  man- 
ganate.  Further,  a  salt  of  the  composition  KH3Mn4OJO 
is  formed  as  a  brown  insoluble  powder  by  boiling  the 
other  manganites  with  potassium  hydroxide  or  car- 
bonate. 

Weldon's  Process  for  the  ^Regeneration  of  Manganese 
Dioxide  in  the  Preparation  of  Chlorine. — Under  the  head 
of  Chlorine  (which  see),  Weldon's  process  was  referred 
to  ;  but  as  a  satisfactory  explanation  of  the  working  of  the 
process  could  not  be  given  without  dealing  with  some 
rather  complicated  compounds  of  manganese,  a  fuller 
account  was  postponed  until  these  compounds  should 
be  taken  up.  The  object  in  view  is  to  utilize  the  waste 
liquors  from  the  chlorine  factories.  When  manganese 
dioxide  is  treated  with  hydrochloric  acid,  as  we  have  seen, 
manganous  chloride  and  chlorine  are  formed,  according 
to  the  equation 

Mn02  +  4HC1  =  MnCl2  +  C12  +  2H2O. 


684  INORGANIC  CHEMISTRY. 

The  manganous  chloride  was  of  no  special  value  until  it 
was  shown  that  by  a  comparatively  simple  method  it  can 
be  converted  into  a  compound  which  with  hydrochloric 
acid  gives  chlorine.  When  it  is  treated  in  solution  with 
lime  the  corresponding  hydroxide  is  precipitated  : 

MnCl,  +  Ca(OH)2  =  Mn(OH)a  +  CaCla  ; 

and  when  this  hydroxide  mixed  with  lime  is  allowed  to 
stand  exposed  to  the  air  oxidation  takes  place,  and  a 
compound  CaMn03  or  CaMn2O6  is  formed  : 

Mn(OH)2    +  Ca(OH)2  +  O    =  CaMnO3  +  2H2O  ; 
2Mn(OH)2  +  Ca(OH)2  +  2O  =  CaMn2O5  +  3H2O. 

These  -compounds  give  chlorine  when  treated  with  hydro- 
chloric acid.  They  may  indeed  be  regarded  as  consisting 
of  lime  and  manganese  dioxide,  CaO.  MnO2  and  CaO.2MnO2, 
and  the  action  of  hydrochloric  acid  takes  place  thus  : 


CaO.Mn02    +  6HC1    =CaCl2  +  MnCl2   +3H2O  +  C12; 
Ca0.2Mn02  +  10HC1  =  CaCl2  +  2MnCl2  +  5H2O  +  2C13. 

In  practice,  the  waste  liquor  is  mixed  with  calcium 
carbonate  in  order  to  neutralize  the  acid.  After  settling, 
lime  enough  is  added  to  precipitate  the  manganese  as 
hydroxide,  and  to  form  with  this  a  mixture  in  molecular 
proportions.  Into  this  mixture  steam  and  air  are  passed, 
when  the  oxidation  referred  to  takes  place,  and  calcium 
nianganite  is  formed. 

Sulphides.  —  When  a  solution  of  a  manganous  salt  is 
treated  with  ammonium  sulphide,  a  flesh-colored  pre- 
cipitate, which  is  thought  to  be  the  hydrosulphide,  is 
formed.  When  this  is  exposed  to  the  air  it  turns  dark 
in  consequence  of  oxidation  ;  and  if  allowed  to  stand  in 
the  liquid,  if  this  is  concentrated,  it  turns  green  and  be- 
comes crystalline.  The  product  thus  formed  is  man- 
ganous sulphide,  MnS.  This  also  occurs  in  nature  as 
alabandite.  A  disulphide,  MnS2,  corresponding  to 
the  dioxide  is  also  found  in  nature,  and  is  known  as 
hauerite. 


VARIOUS  COMPOUNDS  OF  MANGANESE.  685 

Manganous  Cyanide,  Mn(C!N")2,  in  combination  with,  po- 
tassium or  sodium  cyanide  as  the  compounds  Mn(CN)2. 
4KCN  or  K4Mn(CN)6,  and  Mn(CN)2.4NaCN  or  Na4Mn(CN)6, 
is  formed  by  treating  solutions  of  manganous  salts  with 
potassium  or  sodium  cyanides.  When  exposed  to  the 
air,  or  when  the  solutions  are  boiled,  salts  of  the  formulas 
Mn(CN)3.3KCN  or  K3Mn(CN)6  and  Mn(CN)3.3NaCN  or 
Na3Mn(CN)6  are  formed. 

Manganous  Carbonate,  MnCO3,  is  found  in  nature,  and 
is  precipitated  when  a  solution  of  a  manganous  salt  is 
treated  with  a  soluble  carbonate. 

Manganous  Sulphate,  MnSO4,  is  formed  by  heating  the 
oxides  of  manganese  with  concentrated  sulphuric  acid. 
If  higher  oxides  than  manganous  oxide  are  used,  oxygen 
is  given  off : 

MnO    +H2S04    =MnS04    +  H2O ; 

Mn  A  +  2H2S04  =  2MnS04  +  2H2O  +  O  ; 

Mn02  +  H2S04    =  MnSO4    +  H2O    +  O. 

It  crystallizes  at  low  temperatures  with  seven  molecules 
of  water,  and  at  ordinary  temperatures  with  five,  in  this 
respect  resembling  cupric  sulphate  (which  see).  *  The 
salt  of  the  formula  MnSO4  +  7H2O  forms  bright-red 
monoclinic  prisms  ;  while  that  of  the  formula  MnSO4  -f- 
5H2O  forms  pink  triclinic  crystals.  Between  20°  and  30° 
it  forms  monoclinic  prisms  with  four  molecules  of  water. 

Manganic  Sulphate,  Mn2(SO4)3,  is  formed  when  the  oxide 
Mn3O4,  or  the  finely  divided  precipitated  dioxide,  MnO2, 
is  treated  with  sulphuric  acid  at  not  too  high  a  tempera- 
ture. It  forms  a  dark  green  amorphous  powder,  which 
is  easily  decomposed  by  heat  and  by  water.  With  the 
sulphates  of  the  alkali  metals  it  forms  salts  analogous 
to  the  alums,  as  KMn(SO4)2  +  12H2O  and  NH4Mn(SO4)2  + 
12H2O,  in  which  manganese  takes  the  place  of  aluminium. 
This  fact  makes  it  appear  probable  that  in  the  manganic 
compounds  manganese  is  trivalent,  as  aluminium  prob- 
ably is  in  its  compounds. 

Manganic  Acid  and  the  Manganates. — When  an  oxide 
of  manganese  is  treated  with  an  energetic  oxidizing  agent 
in  the  presence  of  a  strong  base  it  is  converted  into  a 


686  INORGANIC  CHEMISTRY. 

manganate,  just  as  the  oxides  of  chromium  are  converted 
into  chromates  and  the  compounds  of  sulphur  into  sul- 
phates. These  three  classes  of  compounds  are  analo- 
gous as  far  as  the  composition  is  concerned,  as  shown  by 
the  formulas 

M2MnO4  M2CrO4  M2SO4 

Manganate  Chromate  Sulphate 

The  manganates  are,  however,  quite  unstable  except  in 
alkaline  solution,  and  when  they  decompose  they  form 
the  permanganates. — Potassium  manganate,  K2MnO4,  is 
formed  by  fusing  manganese  dioxide  with  potassium  hy- 
droxide, when,  if  the  air  is  not  in  contact  with  the  mass, 
the  reaction  takes  place  as  represented  in  the  equation 

3MnO3  +  2KOH  =  K2MnO4  +  Mn2O3  +  H,O. 

It  is  also  made  by  fusing  the  dioxide  with  potassium  hy- 
droxide and  potassium  chlorate,  when  this  reaction  takes- 
place : 

3Mn02  +  6KOH  +  KC1O3  =  3K2MnO4  +  KG1  +  3H,(X 

"When  the  mass  obtained  in  either  way  is  treated  with 
water  a  dark-green  solution  of  the  manganate  is  formed, 
and  by  allowing  this  to  evaporate  at  the  ordinary  tem- 
perature in  a  partial  vacuum,  or  in  an  atmosphere  free 
of  oxygen,  the  salt  is  obtained  in  small  crystals,  which 
are  almost  black.  When  a  solution  of  a  manganate  i& 
treated  with  an  acid,  the  manganic  acid  is  at  once  decom- 
posed into  permanganic  acid  and  manganese  dioxide : 

3H2MnO4  =  2HMn04  +  MnO2  +  2H2O. 

The  change  of  a  manganate  to  a  permanganate  is  ef- 
fected simply  by  passing  carbon  dioxide  into  the  solu- 
tion, or  by  boiling  or  allowing  the  solution  to  stand  in 
the  air.  The  change  by  means  of  carbon  dioxide  is  rep- 
resented by  the  equation 

3KaMnO4  +  2C03  =  2KaC03  +  MnO2  +  2KMnO4. 


POTASSIUM  PERMANGANATE.  687 

With  water  the  change  takes  place  thus : 

3K2MnO4  +  2HaO  =  2KMnO4  +  MnOa  +  4KOH. 

The  potassium  hydroxide  and  the  manganese  dioxide 
react  upon  each  other  to  form  a  manganite  of  more  or 
less  complicated  composition.  While  the  manganates 
are  decomposed  by  acids,  forming  permanganates,  the 
latter  are  decomposed  by  alkalies,  forming  manganates. 
Thus,  when  a  solution  of  potassium  permanganate  is 
boiled  with  potassium  hydroxide  the  color  changes  to 
green,  owing  to  the  formation  of  the  manganate  : 

2KOH  +  2KMnO4  =  2K2MnO4  +  H2O  +  O. 

This  change  takes  place  readily  in  the  presence  of  sub- 
stances which  have  the  power  to  take  up  oxygen ;  but  if 
such  substances  are  present  the  reduction  goes  further, 
forming  finally  a  manganite  which  is  a  derivative  of  the 
hydroxide,  MnO(OH)2. 

Permanganic  Acid  and  the  Permanganates. — The  sim- 
plest method  of  obtaining  the  permanganates  is  by  de- 
composition of  the  manganates,  as  described  in  the  last 
paragraph. 

Potassium  Permanganate,  KMnO4,  is  manufactured  on 
the  large  scale  by  oxidizing  manganese  dioxide  in  the 
presence  of  a  base.  Sometimes  the  oxidation  is  effected 
by  the  oxygen  of  the  air ;  sometimes  by  the  action  of  an 
oxidizing  agent,  as  potassium  chlorate  or  nitrate.  The 
fundamental  reaction  in  each  case  is  that  represented  by 
the  equation 

MnO2  +  2KOH  +  O  =  K2MnO4  +  H2O. 

As  will  be  observed,  it  is  a  reaction  of  the  same  kind  as 
that  involved  in  the  conversion  of  a  sulphite  into  a  sul- 
phate. Probably  the  first  action  of  the  hydroxide  upon 
the  dioxide  consists  in  the  formation  of  the  manganite, 
K2MnO3,  and  this  is  then  oxidized  to  the  manganate. 
When  the  solution  of  the  manganate  is  treated  with  sul- 
phuric acid  a  change  similar  to  those  referred  to  above 
takes  place,  and  the  permanganate  is  formed.  The  salt 


688  INORGANIC  CHEMISTRY. 

is  easily  soluble  in  water,  and  is  deposited  from  its  solu- 
tion in  crystals,  isomorphous  with  potassium  perchlorate, 
which  appear  nearly  black,  with  a  greenish  lustre.  Its- 
solution  in  water  has  a  purple  or  reddish-purple  color, 
according  to  the  concentration.  Yery  concentrated  solu- 
tions appear  almost  black.  The  salt  is  used  extensively 
in  the  laboratory  and  in  the  arts  as  an  oxidizing  agent. 
Its  action  will  be  readily  understood  from  what  has- 
already  been  said  in  regard  to  the  conduct  of  manganese 
towards  acids  and  towards  alkalies.  When  the  perman- 
ganate undergoes  decomposition  in  the  presence  of  an 
acid  the  manganese  tends  to  form  a  manganous  salt,  and 
all  the  oxygen  present  in  excess  of  what  is  needed  for 
this  purpose  is  given  off.  Thus  the  decomposition  with 
sulphuric  acid  takes  place  as  represented  in  the  equation 

2KMn04  +  3H2S04  =  2MnSO4  +  K2SO4  +  3H2O  +  5O. 

Therefore,  when  potassium  permanganate  is  used  as  an 
oxidizing  agent  in  acid  solution,  two  molecules  of  the 
salt  KMnO4  give  five  atoms  of  oxygen.  On  the  other 
hand,  when  the  action  takes  place  in  alkaline  solution 
the  action  reaches  its  limit  in  a  manganite,  which,  for 
purposes  of  calculation,  may  be  regarded  as  having  the 
composition  K2MnO3.  The  first  change  is  from  the  per- 
manganate to  the  manganate  as  represented  in  the  equa- 
tion 

2KMn04  +  2KOH  =  2K2MnO4  +  H2O  +  O, 
and  then  the  manganate  loses  another  atom  of  oxygen, 
2K2MnO4  =  2K2MnO3  +  2O. 

Therefore,  when  the  permanganate  is  used  as  an  oxid- 
izing agent  in  alkaline  solution,  two  molecules  of  the 
salt  yield  three  atoms  of  oxygen. 

The  permanganates  and  maganates  are  valuable  dis- 
infecting agents,  and  the  sodium  salts  are  extensively 
used  for  this  purpose,  under  the  name  of  Candy's  liquid. 

When  a  solution  of  barium  permanganate  is  treated 
with  sulphuric  acid,  free  permanganic  acid  is  obtained 


REACTIONS  OF  MANGANESE  COMPOUNDS.         689 

in  solution.  It  is  extremely  unstable,  and  decomposes 
spontaneously  when  the  solution  is  exposed  to  the  light 
or  is  heated.  When  dry  potassium  permanganate  is 
added  to  concentrated  sulphuric  acid,  oily  drops  sepa- 
rate and  collect  upon  the  bottom  of  the  vessel.  These 
are  manganese  heptoxide,  Mn2O7,  which  is  formed  thus : 

2KMnO4  +  H2SO4  =  K2SO4  +  Mn2O7  +  H2O. 

The  compound  bears  to  permanganic  acid  the  relation 
of  an  anhydride : 

2HMnO4  =  Mn2O7  +  H,O. 

It  is  extremely  unstable,  giving  off  oxygen  with  great 
ease,  and  therefore  acting  as  a  powerful  oxidizing  agent. 
As  regards  the  constitution  of   the  manganates  and 
permanganates,  they  are  respectively  regarded  as  anal- 
ogous   to    the    sulphates    and  perchlorates.      Accord- 
ingly   manganic    acid   is   represented   by   the    formula 
O 

HO-Mn-OH  or  MnO2(OH)2,  while  permanganic  acid  is 

6 

o 

represented  by  the  formula  O=Mn-OH  or  MnO3(OH). 

b 

Reactions  which  are  of  Special  Value  in  Chemical 
Analysis. — The  conduct  of  manganous  salts  towards 
soluble  hydroxides  and  towards  soluble  carbonates  has 
been  described.  The  hydroxide  is  soluble  in  ammonia 
and  ammonium  salts,  but  this  solution  turns  brown  when 
exposed  to  the  air  and  the  manganese  is  gradually  pre- 
cipitated as  the  hydroxide  Mn(OH)3. 

The  conduct  towards  ammonium  sulphide  has  been 
described.  When  oxidizing  agents  like  hypochlorites, 
chlorine,  or  bromine  act  upon  manganous  salts  in  solu- 
tions in  presence  of  soluble  hydroxides,  hydroxides  cor- 
responding to  the  dioxide  MnO2,  such  as  Mn(OH)4, 
MnO(OH)2,  are  precipitated.  Instead  of  the  above- 


690  INORGANIC  CHEMISTRY. 

mentioned  oxidizing  agents,  potassium  permanganate  may 
be  used. 

The  action  of  potassium  permanganate  and  manganate 
as  oxidizing  agents  when  used  in  alkaline  and  in  acid 
solutions  has  been  described  above. 

Manganese  is  easily  detected  by  heating  the  substance 
under  examination  with  nitric  acid  and  lead  peroxide, 
when  permanganic  acid  will  be  formed  if  manganese  is 
present,  and  its  formation  will  be  shown  by  the  purple 
color  of  the  solution. 

"With  microcosmic  salt  and  borax  manganese  gives  an 
amethyst-colored  bead  in  the  oxidizing  flame,  which 
becomes  colorless  in  the  reducing  flame. 


CHAPTER  XXXIII. 

ELEMENTS  OF  FAMILY 'VIII,  SUB-GROUP  A: 
IRON— COBALT— NICKEL. 

General. — The  three  elements  which  form  this  group 
are  in  many  respects  very  similar,  and  their  atomic 
weights  differ  but  little  from  one  another.  That  of  iron 
(55.6)  is  nearly  the  same  as  that  of  manganese  (54.57), 
while  cobalt  and  nickel  have  nearly  the  same  atomic 
weight.  There  is  much  in  iron  which  suggests  manganese. 
It  forms  two  series  of  compounds,  the  ferrous  and  ferric 
compounds,  which  are  analogous  to  the  manganous  and 
manganic  compounds.  In  the  first  series  iron  appears  to 
be  bivalent,  as  shown  in  the  formulas 

FeCla,  Fe(OH)2,  FeO,  FeS,  FeSO4,  FeCO3,  etc. 

In  the  second  series  it  appears  to  be  trivalent,  as  indi- 
cated in  the  formulas 

FeCls,  Fe(OH)3)  Fe,0s,  Fe(NO,)s,  Fe,(SO4)a)  etc. 

Like  chromium  and  manganese  it  also  forms  an  acid 
known  as  ferric  acid,  H2FeO4,  which  in  composition  is 
analogous  to  chromic  and  manganic  acids.  The  soluble 
salts  of  this  acid  are,  however,  unstable,  and  on  decom- 
posing yield  ferric  hydroxide.  Oxidizing  agents  readily 
convert  ferrous  compounds  into  ferric  compounds,  and 
reducing  agents  reconvert  the  latter  into  the  former. 
When  exposed  to  the  air  most  ferrous  compounds  are 
oxidized  to  ferric  compounds.  The  ferrous  compounds 
in  which  iron  is  bivalent  are  similar  to  the  compounds 
of  the  zinc  group.  The  ferric  compounds,  however,  in 
which  the  iron  is  trivalent,  are  similar  to  the  aluminium 
compounds  ;  and  in  ferric  acid  it  exhibits  a  resemblance 
to  chromium.  Cobalt  and  nickel  resemble  iron  in  re- 

(691) 


692  INORGANIC  CHEMISTRY. 

spect  to  their  power  to  form  two  series  of  compounds 
corresponding  to  the  ferrous  and  ferric  compounds. 
Both  elements  preferably  form  compounds  of  the  lower 
series,  examples  of  which  are  represented  by  the  for- 
mulas 

CoCla      Co(OH)2      CoO      Co(NO3)2      CoSO4      etc. 
MCI,      Ni(OH)2      MO      Ni(N03)a      NiSO4      etc. 

Cobalt  forms  a  few  compounds  corresponding  to  the 
ferric  series ;  and  nickel  forms  a  hydroxide,  of  the  for- 
mula Ni(OH)3.  While  the  power  of  cobalt  to  form  com- 
pounds in  which  it  is  trivalent  is  much  weaker  than  that 
of  iron,  it  is  stronger  than  that  of  nickel,  the  latter  being 
almost  exclusively  bivalent.  In  general  terms,  it  may  be 
said  that  manganese  forms  a  greater  variety  of  compounds 
than  any  other  element  except  carbon.  In  the  manganous 
compounds  it  exhibits  analogies  with  zinc,  copper,  and 
some  other  bivalent  elements;  in  the  manganic  compounds 
it  exhibits  analogies  with  aluminium  ;  in  manganic  acid  it 
suggests  sulphur  and  chromium  ;  and  in  permanganic 
acid  it  suggests  chlorine.  In  the  following  table  some 
of  the  analogies  which  are  plainly  discernible  between 
the  elements  mentioned,  and  iron,  cobalt,  and  nickel,  are 
indicated.  The  formulas  of  those  compounds  which  are 
not  easily  obtained,  and  are  exceptional,  are  put  in 
brackets : 

MnSO4         [Mn2(SO4)3]        MnOa       K2MnO4        KMnO4 

[CrS04]        Cra(S04)a  CrO3         K2CrO4          [HCrO4](?) 

FeSO4          Fe2(SO4)8  [K2FeO4] 

CoSO4          Co(OH)3 
NiS04          [Ni(OH)3] 
A12(S04)3 
ZnSO4 

SO2          SO3          K2SO4 

KC104 

As  regards  the  question  whether  the  formula  of  the 
simpler  ferric  compounds  is  to  be  written  with  two  atoms 
of  iron  in  every  case,  it  is  in  much  the  same  state  as  the 
question  in  regard  to  aluminic  compounds.  Is  ferric 
chloride  FeCl3  or  Fe2Cl6?  A  determination  of  the 


VALENCE  OF  IRON.  693 

specific  gravity  of  the  vapor  gave  a  result  in  accordance 
with  the  larger  formula,  but  this  would  not  appear  to  be 
sufficient  evidence  in  view  of  the  peculiar  results  ob- 
tained with  aluminium  chloride.  Considering  the  close 
resemblance  between  ferric  compounds  and  the  com- 
pounds of  aluminium,  it  seems  probable  that,  if  alu- 
minium is  trivalent,  iron  is  also  trivalent  in  these  com- 
pounds. The  ease  with  which  ferric  chloride  forms  com- 
pounds with  other  chlorides  suggests,  further,  that  the 
compound  of  the  simpler  formula  FeCl8  may  combine 
with  another  molecule  of  the  same  kind  to  form  a  double 

/(Cl,)\ 
chloride  of  the  formula  Fe,Cl.  or  Fef-(Cl,HFe.     It  may 


be  objected  to  this  that  it  is  not  probable  that  such  a 
compound  could  be  converted  into  vapor  without  under- 
going decomposition,  and,  according  to  the  one  determi- 
nation of  the  specific  gravity,  it  does  not  appear  to 
undergo  decomposition.  What  value  to  attach  to  this 
objection  it  is  impossible  to  say  at  present.  In  any  case, 
a  further  knowledge  of  the  facts  is  needed  before  a  final 
conclusion  can  be  reached.  In  the  mean  time  it  seems 
to  be  justifiable  to  consider  iron  trivalent  in  ferric  com- 
pounds, as  aluminium  is  considered  trivalent  in  its  com- 
pounds, chromium  in  chromic  compounds,  and  manga- 
nese in  manganic  compounds. 

That  iron  is  bivalent  in  ferrous  compounds  is  probable 
from  the  analogy  of  these  compounds  with  the  distinctly 
bivalent  metals,  like  copper,  zinc,  etc.  Further,  a  deter- 
mination of  the  specific  gravity  of  the  vapor  of  ferrous 
chloride  gave  a  figure  which  indicated  that  the  vapor 
consisted  of  about  an  equal  number  of  molecules  of  the 
formulas  FeCl2  and  Fe2Cl4,  so  that  it  appears  that  at  a 
lower  temperature  the  compound  has  the  formula  Fe2Cl4, 
and  that  the  compound  breaks  down  or  dissociates,  form- 
ing the  simpler  compound.  This  subject  requires  further 
investigation.  In  the  mean  time  the  simpler  formula  will 
be  used,  as  it  probably  represents  the  chemical  molecule 
or  that  smallest  particle  of  the  compound  which  comes 
into  play  in  chemical  reactions.  If  iron  is  bivalent  in 


694  INORGANIC  CHEMISTRT. 

ferrous  compounds,  then  in  all  probability  cobalt  and 
nickel  are  bivalent  in  their  principal  compounds. 

IRON,  Fe  (At.  Wt.  55.6). 

Introductory. — The  importance  of  this  metal  to  man- 
kind can  hardly  be  overestimated,  and  for  many  cen- 
turies it  has  played  a  commanding  part  in  the  industries. 
It  requires  little  thought  to  convince  one  that  without  it 
the  earth  would  be  quite  a  different  place  from  what  it 
now  is.  In  the  earliest  periods  of  history  metals  were 
but  little  used,  as  but  few  of  them  are  furnished  ready 
for  use  by  nature.  Stones  were  therefore  first  used,  and 
these  were  shaped  into  a  variety  of  implements,  many  of 
which  still  exist,  and  furnish  evidence  of  the  Stone  Age. 
After  a  time  copper  and  tin  were  used  in  the  form  of  an 
alloy  or  bronze,  as  copper  is  found  in  nature  in  the  free 
condition.  During  this  period,  known  as  the  Bronze  Age, 
stone  implements  gave  way  to  those  made  of  bronze. 
Afterwards  men  learned  to  extract  iron  from  its  ores, 
and  the  Iron  Age  was  introduced  ;  and  this  has  continued 
up  to  the  present,  as  nothing  has  since  been  found  which 
can  advantageously  take  the  place  of  iron.  The  sugges- 
tion has  been  made  that  as  it  is  less  difficult  to  extract 
iron  from  its  ores  than  to  make  bronze,  possibly  iron 
was  used  as  early  as  bronze — perhaps  earlier,  but  that, 
owing  to  the  fact  that  iron  easily  rusts,  implements  of 
this  metal  have  disappeared,  while  those  made  of  bronze 
remain  intact. 

Forms  in  which.  Iron  occurs  in  Nature. — Iron  occurs  in 
small  quantity  native  in  meteorites,  in  the  basalts  of  Bo- 
hemia and  Greenland,  and  in  some  gabbros.  The  iron 
meteorites  always  contain  nickel,  and  frequently  small 
quantities  of  other  elements,  as  manganese  and  carbon. 
Compounds  of  iron  occur  in  enormous  quantities,  and 
widely  distributed  in  the  earth.  Among  the  more  im- 
portant are  the  following-named :  hematite,  Fe2O3 ;  mag- 
netite, Fe3O4 ;  brown  iron  ore,  Fe4O3(OH)6 ;  siderite,  or 
the  carbonate,  FeCO3 ;  pyrite,  FeS2 ;  pyrrhotite,  Fe7S8. 
It  is  also  contained  in  many  silicates  in  small  quan- 
tity, and  in  consequence  of  the  disintegration  of  the 


METALLURGY  OF  IRON.  695 

constituents  of  rocks  it  is  found  in  the  soil,  and  in  many 
natural  waters.  In  the  vegetable  kingdom  it  is  always 
found  in  chlorophyll,  and  in  the  animal  kingdom  always 
in  the  blood.  The  compounds  which  are  chiefly  used 
for  the  purpose  of  making  iron,  or  the  iron  ores,  are 
magnetite,  Fe3O4 ;  hematite,  Fe2O3 ;  brown  iron  ore, 
Fe4O3(OH)6 ;  and  spathic  iron,  or  siderite,  FeCO3. 

Metallurgy. — The  ores  of  iron,  after  they  are  broken 
up,  are  first  roasted,  in  order  to  drive  off  water  from  the 
hydroxides ;  to  decompose  carbonates ;  to  oxidize  sul- 
phides ;  and,  as  far  as  possible,  to  convert  the  oxides 
into  ferric  oxide,  Fe2O3,  which  is  the  most  easily  re- 
ducible of  the  oxides  of  iron.  After  the  ores  are  pre- 
pared in  this  way  they  are  reduced  by  heating  them  with 
carbon  and  fluxes  in  the  blast- 
furnaces, when  the  iron  collects 
in  the  molten  condition  under 
the  so-called  slag  at  the  bottom 
of  the  furnace.  Blast-furnaces 
differ  somewhat  in  construction, 
but  the  essential  parts  are  rep- 
resented in  Fig.  14. 

The  inner  cavity  of  the  furnace 
is  narrow  at  the  top  and  bot- 
tom, as  is  shown  in  the  fig- 
ure. Through  pipes,  known  as 
tuyeres,  such  as  that  represented 
at  the  lower  part  of  the  left- 
hand  side  of  the  figure,  air  is 
blown  into  the  furnace  to  facili- 
tate the  combustion.  In  modern 
furnaces  arrangements  are  made 
above  for  carrying  off  the  gases  Fl°- «. 

and  utilizing  them  as  fuel.  The  inner  walls  are  built 
of  fire-bricks,  and  these  are  surrounded  by  ordinary 
bricks,  or  stone-work.  The  furnaces  vary  in  height 
from  25  to  80  or  90  feet,  an  average  height  being  about 
45  feet.  The  reduction  of  the  ores  is  accomplished  bv 
placing  in  the  furnace  alternating  layers  of  coke  or 
charcoal,  and  the  ores  mixed  with  proper  fluxes.  The 


696  INORGANIC  CHEMISTRY. 

nature  of  the  flux  depends  upon  the  ore.  If  this  con- 
tains silicon  dioxide  or  clay,  lime  is  added ;  while, 
if  it  contains  considerable  lime,  minerals  rich  in  silicic 
acid  are  used,  such  as  feldspar,  clay-slate,  etc.  The 
object  of  the  flux  is  to  form  a  slag  in  which  the  re- 
duced iron  collects,  and  by  which  it  is  protected  from 
oxidation.  When  the  fire  is  once  started  in  a  blast- 
furnace the  operation  of  reduction  is  continuous  until 
the  furnace  is  burned  out.  Alternate  layers  of  ore  and 
flux  and  carbon  are  added,  and,  as  the  reduced  iron  col- 
lects below,  it  is  from  time  to  time  drawn  off  and  allowed 
to  solidify  in  moulds  of  sand.  The  operation  requires 
close  attention.  The  ores  must  be  carefully  studied,  and 
the  nature  and  amount  of  flux  regulated  according 
to  the  character  of  the  ore  as  above  stated.  Then, 
too,  the  temperature  of  the  furnace  is  a  matter  of  im- 
portance, and  must  be  watched,  and  regulated  by  means 
of  the  blast.  The  reduction  is  largely  accomplished 
by  carbon  monoxide.  In  the  lower  part  of  the  furnace 
the  fuel  burns  to  carbon  dioxide,  but  this  comes  in  con- 
tact with  hot  carbon,  and  is  then  reduced  to  the  monox- 
ide. The  hot  monoxide  in  contact  with  the  oxides  of 
iron  reduces  these,  and  is  itself  converted  into  the  diox- 
ide. A  large  proportion  of  the  carbon  monoxide,  how- 
ever, escapes  oxidation,  and  this  is  carried  off  from  the 
top  of  the  furnace  to  the  bottom  by  properly  arranged 
pipes,  and  is  then  utilized  as  fuel.  A  furnace  lasts  from 
two  to  twenty  years,  and  sometimes  longer. 

Varieties  of  Iron. — The  iron  obtained  as  above  de- 
scribed is  known  as  pig-iron  or  cast-iron.  It  is  very 
impure,  containing  carbon,  phosphorus,  sulphur,  silicon, 
etc.  If,  when  drawn  from  the  furnace,  the  iron  is  cooled 
rapidly,  nearly  all  the  carbon  contained  in  it  remains  in 
chemical  combination,  and  the  iron  has  a  silver-white 
color.  This  product  is  known  as  white  cast-iron.  If  the 
iron  cools  slowly,  most  of  the  carbon  separates  as  graph- 
ite, and  this  being  distributed  through  the  mass  gives  it 
a  gray  color.  This  product  is  known  as  gray  cast-iron. 
If  the  ore  contains  considerable  manganese,  this  is  re- 
duced with  the  iron,  and  iron  made  from  such  ores  and 


VARIETIES  OF  IRON.  697 

containing  manganese  has  the  power  to  take  up  more 
carbon  than  ordinary  iron.  This  product,  containing 
from  3.5  to  6  per  cent  combined  carbon,  is  known  as 
spiegel-iron. 

All  varieties  of  cast-iron  are  brittle,  and  easily  fusible. 
The  gray  iron  fuses  at  a  lower  temperature  than  the 
white,  and  is  not  as  brittle  ;  it  is  therefore  well  adapted  to 
making  castings.  When  cast-iron  is  treated  with  hydro- 
chloric acid  the  carbon  which  is  present  in  combined 
form  is  given  off  in  combination  with  hydrogen  as  hy- 
drocarbons, some  of  which  have  a  disagreeable  odor. 
This  is,  of  course,  the  cause  of  the  bad  odor  noticed  in 
dissolving  ordinary  cast-iron  in  acids.  The  uncombined 
or  graphitic  carbon,  on  the  other  hand,  remains  undis- 
solved.  Owing  to  its  brittleness,  cast-iron  cannot  be 
welded.  When  the  carbon,  silicon,  and  phosphorus 
are  removed  the  iron  becomes  tough  and  malleable,  and 
its  melting-point  is  much  raised.  The  product  thus  ob- 
tained is  known  as  wrought-iron. 

Puddling. — Wrought-iron  is  obtained  from  cast-iron  by 
the  puddling  process.  The  puddling  furnace  has  a  flat, 
oval  hearth,  and  low  arched  roof.  The  sides  of  the 
hearth  are  lined  with  a  layer  of  iron  ore  (oxide).  Coal 
is  burned  on  a  grate  and  the  flame  passes  into  the  fur- 
nace at  one  end  and  out  at  the  other,  thus  coming  in 
contact  with  the  roof  and  the  charge  of  iron.  By  con- 
tact with  the  flame,  and  by  the  heat  radiated  from  the 
roof,  the  cast-iron  melts.  The  carbon  and  silicon  are 
removed  from  the  molten  cast-iron,  partly  by  the  oxy- 
gen in  the  air  or  flame,  but  principally  by  the  oxygen 
in  the  iron  ore,  which  is  itself  thus  reduced  to  wrought- 
iron. 

Wrought-iron  contains  less  than  0.6  per  cent  of  car- 
bon, and,  as  the  percentage  of  carbon  decreases,  the. 
malleability  increases  and  the  melting-point  rises.  The 
melting-point  of  good  wrought-iron  is  from  1900°  to 
2100°.  Small  quantities  of  sulphur,  phosphorus,  silicon, 
and  manganese  exert  a  very  marked  influence  upoji  its 
properties.  The  process  of  welding  consists  in  heating 


698  INORGANIC  CHEMISTRY. 

two  pieces  of  iron  to  a  high  temperature,  putting  some 
borax  upon  one  of  them,  laying  them  together,  and  ham- 
mering, when,  as  is  well  known,  they  adhere  firmly  to- 
gether. The  object  of  the  borax  is  to  keep  the  surfaces 
bright,  which  it  does  by  uniting  with  the  oxide  and  form- 
ing an  easily  fusible  borate. 

Bessemer  Process. — Molten  cast-iron  is  poured  into  a 
large  vessel  called  the  converter.  The  carbon  and  sili- 
con are  entirely  oxidized  and  removed  by  means  of  a 
blast  of  air  forced  through  the  metal  from  below.  No 
fuel  is  used,  as  the  heat  generated  by  the  oxidation  of 
carbon  and  silicon  is  sufficient  to  raise  the  temperature 
above  2100°.  The  converter  contains  molten  wrought- 
iron  after  the  oxidation.  By  addition  of  spiegel-iron  & 
product  containing  any  desired  percentage  of  carbon  is 
obtained. 

Iron  which  contains  more  than  a  very  small  percent- 
age of  phosphorus  is  not  adapted  to  the  manufacture  of 
Bessemer  steel  in  the  ordinary  way  ;  but  it  has  been 
found  that,  if  the  converters  are  lined  with  lime  and 
magnesia,  such  iron  may  be  used.  Under  these  circum- 
stances the  phosphorus  is  oxidized,  and  forms  calcium 
and  magnesium  phosphates,  which  are  of  value  as  fertil- 
izers (see  Calcium  Phosphate).  This  process  is  known 
as  the  Thomas-Gilchrist  or  the  basic-lining  process. 

Siemens- Martin  Furnace. — This  is  simply  a  reversible 
puddling  furnace  in  which  gas  is  used  as  fuel.  The 
gas  is  previously  heated  in  a  Siemens  regenerative  fur- 
nace. 

Steel  and  Wrought-iron. — The  product  of  the  puddling 
furnace  is  called  wrought-iron ;  while  those  formed  in 
the  Bessemer  process  and  in  the  Siemens-Martin  fur- 
nace are  called  steel.  Bessemer  steel  often  contains 
less  than  0.6  per  cent  of  carbon,  and  Siemens-Martin 
steel  is  the  purest  form  of  wrought-iron,  containing  less 
carbon  and  silicon  than  the  product  of  the  puddling 
furnace. 

Tempering. — When  steel  is  heated  and  cooled  sud- 
denly, it  is  rendered  extremely  hard  and  brittle ;  and 
when  hardened  steel  is  carefully  heated,  and  allowed  to- 


PROPERTIES   OF  IRON. 

cool  slowly,  it  becomes  very  elastic.     This  process  is 
called  tempering. 

Properties  of  Iron. — Pure  iron  is  almost  unknown.  Of 
the  commercial  varieties,  it  follows  from  what  has  been 
said  that  wrought-iron  is  the  purest.  That  which  is 
used  for  piano- strings  is  the  purest  iron  which  can  be 
bought ;  it  contains  only  about  0.3  per  cent  of  impuri- 
ties. Pure  iron  can  be  made  in  the  laboratory  by  ignit- 
ing the  oxide  or  oxalate  in  a  current  of  hydrogen,  and 
by  reducing  ferrous  chloride  in  hydrogen.  In  larger 
quantity  it  can  be  prepared  by  melting  the  purest 
wrought-iron  in  a  lime  crucible  by  means  of  the  oxy hy- 
drogen flame.  The  impurities  are  taken  up  by  the  cru- 
cible, and  a  regulus  of  the  pure  metal  is  left  behind. 
That  made  by  reduction  of  the  oxide  or  oxalate  is,  of 
course,  in  finely  divided,  condition.  If  in  its  preparation 
the  temperature  is  kept  as  low  as  possible,  the  prod- 
uct takes  fire  when  brought  in  contact  with  the  air ; 
while  if  the  temperature  is  high,  the  product  has  not 
this  power.  Iron  is  white,  and  is  one  of  the  hardest 
metals ;  and  its  melting-point  is  higher  than  that  of 
wrought-iron.  Pure  iron  is  attracted  by  the  magnet. 
In  contact  with  a  magnet,  or  when  placed  in  a  coil 
through  which  an  electric  current  is  passing,  it  becomes 
a  magnet ;  but  the  purer  it  is  the  sooner  it  loses  the  mag- 
netic power  when  removed  from  the  magnet  or  the  coil. 
Steel,  however,  retains  its  magnetism.  When  heated  to 
a  sufficiently  high  temperature  iron  burns,  and  forms 
the  oxide,  Fe3O4.  This  takes  place  much  more  easily  in 
oxygen  than  in  the  air.  In  dry  air  iron  does  not  under- 
go change,  but  in  moist  air  it  rusts,  or  it  becomes  covered 
with  a  layer  of  oxide  and  hydroxide,  which  is  formed 
by  the  action  of  the  air,  carbon  dioxide,  and  water. 
Water  that  contains  salts  in  solution  facilitates  the 
rusting.  Various  methods  are  adopted  to  protect  iron 
from  this  change,  most  of  which  are,  however,  purely 
mechanical.  A  method  which  promises  valuable  results 
is  that  invented  by  Barff,  which  consists  in  introducing 
the  iron  into  water  vapor  at  a  temperature  of  650°,  when 
it  becomes  covered  with  a  firmly  adhering  layer  of  oxide. 


700  INORGANIC  CHEMISTRY. 

Iron  dissolves  in  acids  with  evolution  of  hydrogen, 
and  generally  with  formation  of  ferrous  salts  : 


Fe  +  2HCl   =  FeCl2  +  H2  ; 
Fe  +  H2S04  =  FeS04  +  HQ. 

When  cold  nitric  acid  is  used,  ferrous  nitrate  and  am- 
monium nitrate  are  the  products  ;  if  the  acid  is  warmed, 
ferric  nitrate  and  oxides  of  nitrogen  are  formed.  When 
an  iron  wire  which  has  been  carefully  polished  is  intro- 
duced for  an  instant  into  red  fuming  nitric  acid  it  can 
afterward  be  put  into  ordinary  nitric  acid  without  under- 
going change.  It  is  said  to  be  in  the  passive  state  ;  and 
the  commonly  accepted  explanation  of  the  phenomenon 
is,  that  the  wire  is  covered  with  a  thin  layer  of  oxide. 
As,  however,  the  passive  condition  is  lost  by  contact 
with  an  ordinary  wire,  the  explanation  does  not  appear 
to  be  adequate. 

Ferrous  Chloride,  FeOla.  —  When  iron  is  dissolved  in 
hydrochloric  acid  without  access  of  air,  and  the  solution 
evaporated,  crystals  of  the  composition  FeCl2  -f-  4H2O 
are  obtained.  When  heated  for  the  purpose  of  driving 
off  the  water,  the  crystallized  compound  decomposes. 
The  dry  chloride  can  be  obtained  by  heating  iron  in  a 
current  of  dry  hydrochloric-acid  gas.  It  is  a  colorless 
mass,  which  deliquesces  in  the  air,  is  volatile  at  a  high 
temperature  ;  and  determinations  of  the  specific  gravity 
of  its  vapor  made  at  very  high  temperatures  have  shown 
that  its  molecule  under  these  conditions  should  be 
represented  by  the  formula  FeCl2.  At  lower  tempera- 
tures the  molecule  appears  to  be  more  complex.  The 
evidence  on  this  point  is  not  conclusive.  If  allowed  to 
stand  in  contact  with  the  air  in  hydrochloric-acid  solu- 
tion, it  is  changed  to  ferric  chloride  : 

2FeCl2  +  2HC1  +  O  =  2FeCl3  +  H2O. 

If  hydrochloric  acid  is  not  present,  a  basic  chloride  is 
precipitated  and  ferric  chloride  is  then  in  the  solution  : 


4FeCl,  +  H2O  +  O  =  2Fe  <        +  2FeCl3. 
When  treated  with  oxidizing  agents  in  general,  as  nitric 


FERRIC  CHLORIDE.  701 

acid,  potassium  chlorate,  potassium  permanganate,  etc., 
it  is  converted  into  ferric  chloride. 

Ferrous  chloride,  like  most  other  metallic  chlorides, 
combines  with  the  chlorides  of  the  strongest  basic  ele- 
ments, forming  double  compounds.  Those  with  potas- 
sium and  sodium  chlorides  have  the  formulas  FeCla. 
2KC1  or  K2FeCl4,  and  FeCl2.2NaCl  or  Na2FeCl4.  It 
combines  also  with  other  chlorides,  such  as  those  of 
mercury  and  cadmium,  forming  similar  salts. 

A  solution  of  ferrous  chloride  made  by  dissolving  iron 
in  hydrochloric  acid  is  used  in  medicine  under  the  name 
Liquor  Ferri  chlorati.  It  contains  ten  per  cent  iron. 

Ferric  Chloride,  FeCl3. — As  stated  in  the  last  paragraph, 
ferrous  chloride  is  readily  converted  into  ferric  chloride 
by  oxidation.  The  simplest  way  to  make  a  solution  of 
the  ferric  compound  is  to  dissolve  iron  in  hydrochloric 
acid  and  pass  chlorine  into  it  to  complete  saturation. 
The  solution  is  decomposed  by  heating,  especially  if 
dilute,  yielding  hydrochloric  acid  and  an  insoluble 
oxychloride.  The  chloride  can  be  obtained  in  yellow 
crystals  with  six  or  twelve  molecules  of  water.  Like  the 
ferrous  compound,  it  is  decomposed  into  hydrochloric 
acid  and  the  oxide  when  heated.  Anhydrous  ferric 
chloride  is  obtained  by  heating  iron  wire  in  dry  chlorine. 
It  forms  black,  lustrous  crystalline  laminae,  is  volatile 
&i  a  lower  temperature  than  the  ferrous  compound,  and 
the  specific  gravity  of  the  vapor  is  that  required  by  a 
compound  whose  molecule  corresponds  to  the  formula 
Fe013.  "When  treated  with  nascent  hydrogen,  ferric 
chloride  is  converted  into  ferrous  chloride  : 

FeCls  +  H  =  FeCl2  +  HC1. 

It  combines  with  other  chlorides,  forming  double  chlo- 
rides. A  solution  of  ferric  chloride  is  used  in  medicine 
under  the  name  Liquor  Ferri  sesquichlorati. 

Cyanides. — The  compounds  which  iron  forms  with 
cyanogen  are  of  special  interest.  The  simple  com- 
pounds, ferrous  cyanide,  Fe(CN)2,  and  ferric  cyanide, 
Fe(CN),,  corresponding  to  the  above-mentioned  chlorides, 
are  not  known :  only  double  compounds  of  these  with 


702  INORGANIC  CHEMISTRY. 

other  cyanides  are  well  known,  and  some  of  them  are 
manufactured  on  the  large  scale.  When  a  solution  of 
potassium  cyanide  acts  upon  metallic  iron  or  the 
oxides  of  iron,  a  solution  is  formed  from  which  the 
salt  known  as  potassium  ferrocyanide  or  yelloiv  prus-  - 
slate  of  potash  crystallizes.  This  has  the  composition 
K4Fe(CN)6  -f-  3H2O,  and  may  be  regarded  as  made  up 
of  a  molecule  of  ferrous  cyanide  and  four  molecules 
of  potassium  cyanide,  as  represented  in  the  formula 
Fe(CN)2.4KCN  +  3H2O.  When  this  salt  is  treated  with 
chlorine  it  is  converted  into  potassium  ferricyanide,  or  red 
prussiate  of  potash,  K3Fe(CN)6,  which  is  to  be  regarded 
as  consisting  of  ferric  cyanide  and  potassium  cyanide,  as 
represented  in  the  formula  Fe(CN)3.3KCN.  The  trans- 
formation is  represented  thus : 

K4Fe(CN)6  +  01  =  K3Fe(CN)6  +  KC1. 

From  these  two  a  number  of  other  cyanogen  compounds 
are  obtained.  When  treated  in  concentrated  solution 
with  concentrated  hydrochloric  acid  they  yield  the  free 
acids,  and  by  treating  them  with  solutions  of  different 
metallic  salts  corresponding  salts  of  these  acids  are  ob- 
tained. Among  the  most  important  of  these  derivatives 
are  the  following : 

Ferrohydrocyanic  acid,  Ferrihydrocyanic  acid, 

H4Fe(CN)6  H3Fe(CN)6 

Potassium     ferrocyanide,  Potassium  ferricyanide, 

K4Fe(C]Sr)6  K3Fe(CN)6 

Sodium  ferrocyanide,  Sodium  ferricyanide, 

Na4Fe(ClSr)6  Na3Fe(CN)6 

Barium  ferrocyanide, 

Ba2Fe(CN)« 

Ferric  ferrocyanide,  Ferrous  ferricyanide, 

Fe4[Fe(CN)6]3  Fe3[Fe(CN)6]3 

Ferri- potassium  ferrocyanide, 

KFeFe(CN)8 

Potassium  Ferrocyanide,  K4Pe(CN)6  +  3H2O. — As  stated 
above,  this  salt  can  be  made  by  treating  iron  or  the 
oxides  of  iron  with  a  solution  of  potassium  cyanide.  On 
the  large  scale  it  is  manufactured  by  melting  crude 
potash  or  potassium  carbonate,  and  gradually  adding  a 
mixture  of  iron  filings  or  turnings,  and  refuse  animal- 


CYANIDES  OF  IRON.  703 

matter,  as  claws,  horns,  hoofs,  hair,  etc.  Or  the  potash  is 
melted  with  the  animal  substances  and  potassium  cyanide 
thus  formed,  and  this  treated  in  solution  with  ferrous 
carbonate,  when  the  ferrocyanide  is  formed.  It  forms 
large  yellow  pyramids  belonging  to  the  tetragonal  sys- 
tem. At  the  ordinary  temperature  it  dissolves  in  three 
to  four  parts  of  water,  and  more  easily  in  hot  water.  It 
gives  up  its  water  of  crystallization  very  easily.  When 
heated  it  is  decomposed,  forming  potassium  cyanide, 
nitrogen,  and  a  compound  of  iron  and  carbon : 

K4Fe(CN)6  =  4KCN  +  N2  +  FeC2. 

Treated  with  concentrated  sulphuric  acid  it  undergoes 
decomposition,  giving  as  gaseous  product  carbon  mon- 
oxide, and  this  furnishes  a  good  method  for  the  prep- 
aration of  the  gas : 

K4Fe(CN)6  +  6H2S04  +  6H2O  =  FeSO4  +  2K2S04 

+  3(NH4)2S04  +  6CO. 

With  dilute  sulphuric  acid  it  gives  hydrocyanic  acid,  and 
forms  at  the  same  time  a  white  insoluble  compound  of 
the  composition  KFe(CN)3  or  Fe(CN)2.KCN  : 

2K4Fe(CN)6  +  3H2SO4  =  6HCN  +  3K2SO4  +  2KFe(CN)3. 

Ferrohydrocyanic  Acid,  H4Fe(CN)6,  formed  as  above 
described,  is  a  white  crystalline  substance, which  is  easily 
soluble  in  water  and  alcohol.  It  takes  up  oxygen  from 
the  air,  and  is  converted  into  the  ferric  salt  of  the  acid, 
hydrocyanic  acid  being  given  off.  The  ferric  salt  is 
the  substance  commonly  called  insoluble  Prussian  blue. 
The  relation  of  the  salt  to  the  acid  is  shown  by  the 
formulas 

H.Fe(CN).  Fe,[Fe(CN).], 

Ferro-hydrocyanic  acid  Ferric  ferrocyanide,  or 

Prussian  blue 

Ferric  Ferrocyanide,  or  Prussian  Blue,  Fe4[Fe(C]S")«]3. — 
This  compound  is  very  readily  formed  by  adding  a  solu- 
tion of  a  ferric  salt  to  a  solution  of  potassium  ferrocya- 
nide, and  appears  as  a  dark-blue  precipitate  : 

3K4Fe(CN)6  +  4FeCl3  =  Fe4[Fe(CN)J3  +  12KC1. 


704  INORGANIC  CHEMISTRY. 

It  is  obtained  in  pure  condition  by  treating  a  solution  of 
a  ferric  salt  with  a  solution  of  ferrohydrocyanic  acid. 
When  a  ferric  salt  is  added  to  an  excess  of  potassium 
ferrocyanide  a  ferri-potassium  salt,  KFeFe(CN)6,  is 
formed.  This  is  commonly  called  Prussian  blue,  and 
the  commercial  article  always  contains  some  of  it.  It 
is  also  known  as  soluble  Prussian  blue.  When  heated 
with  an  alkaline  hydroxide,  Prussian  blue  is  decom- 
posed, the  products  of  the  action  being  the  ferrocyanides 
of  the  alkali  metals  and  ferric  hydroxide  : 

Fe4[Fe(CN)6]3  +  12KOH  =  3K4Fe(CN)6  +  4Fe(OH)3. 

Potassium  Ferricyanide,  K3Fe(CN)6.  —  This  salt  is 
formed  by  treating  the  ferrocyanide,  either  dry  or  in 
solution,  with  chlorine.  It  forms  large,  dark-red,  mono- 
clinic  prisms.  It  dissolves  in  about  three  times  its- 
weight  of  water  at  the  ordinary  temperature,  and  is  more 
easily  soluble  in  hot  water.  In  alkaline  solution  it  acts 
as  a  strong  oxidizing  agent,  on  account  of  its  tendency  to 
form  the  ferrocyanide.  The  character  of  the  action  is 
indicated  by  the  following  equation  : 

6K3Fe(CN)6  +  6KOH  =  6K4Fe(CN)6  +  3H2O  +  3O. 

Ferrihydrocyanic  Acid,  H3Fe(CN)6,  is  a  crystallized 
substance. 

Ferrous  Ferricyanide,  Fe3[Fe(CN)6]2,  is  commonly  called 
Turnbull's  blue.  It  is  formed  by  adding  potassium 
ferricyanide  to  a  solution  of  ferrous  sulphate,  or  any 
ferrous  salt : 

3FeS04  +  2K3Fe(CN)6  =  Fe3[Fe(CN)6]2  +  3K2SO4. 

Starting  with  ferrocyanic  and  ferricyanic  acids,  four 
iron  salts  suggest  themselves.  These  are  ferrous  and 
ferric  ferrocyanide,  and  ferrous  and  ferric  ferricyanide. 
The  relations  between  them  are  indicated  in  the  follow- 
ing formulas  : 

Acid H4Fe(CN)6  Acid,    ....       H3Fe(CN)a 

(1)  Ferrous  salt,     .  FeaFe(CN)e  (3)  Ferrous  salt,    .  Fe3[Fe(CN)6]» 

(2)  Ferric  salt,  .     .  Fe4[Fe(CN)6]3        (4)  Ferric  salt,  .     .  FeFe(CN)8 


IRON  SALTS-NITROPRUSSIATES.  705 

Of  these  (2)  is  Prussian  blue  and  (3)  is  Turnbull's 
blue.  The  commercial  Prussian  blue  contains  some 
Turnbull's  blue.  The  reason  of  this  appears  to  be  that 
a  part  of  the  ferrocyanide  of  potassium  used  in  the 
preparation  is  oxidized  by  the  ferric  salt,  and  thus  ferri- 
cyanide  of  potassium  and  a  ferrous  salt  come  together. 
When  potassium  ferrocyanide  is  added  to  a  solution  of 
a  ferrous  salt,  we  should  expect  the  formation  of  salt  (1) 
or  ferrous  ferrocyanide : 

K4Fe(CN)6  +  2FeCl2  =  Fe2Fe(CN)6  +  4KC1. 

But  instead  ol  this  a  ferro-potassium  salt,  of  the  formula 

K2FeFe(CN)6,  is  formed  :' 

K4Fe(CN)6  +    FeCl2  =  K2FeFe(CN)6  +  2KC1. 

This  is  a  white  powder,  which  is  formed  also  when  po- 
tassium ferrocyanide  is  decomposed  by  dilute  sulphuric 
acid  in  the  preparation  of  hydrocyanic  acid.  It  is  repre- 
sented above  by  the  formula  KFe(CN)3,  but  taking  the 
method  of  formation  into  consideration  the  formula 
K2FeFe(CN)6  seems  more  probable.  Ferric  ferricyanide 
is  not  known.  When,  however,  a  solution  of  a  ferric 
salt  is  added  to  one  of  potassium  ferricyanide  the  solu- 
tion turns  dark  brown,  and  perhaps  contains  this  salt. 
The  composition  of  the  salt  is  the  same  as  that  of  ferric 
cyanide,  and  possibly  the  two  compounds  are  identical. 

Nitroprussiates. — When  potassium  ferrocyanide  is 
treated  with  nitric  acid,  potassium  nitrate  is  formed. 
When  this  is  removed  and  the  solution  neutralized  with 
sodium  carbonate,  a  salt  known  as  sodium  nitroprussiate 
is  obtained.  This  crystallizes  very  beautifully,  and  is 
used  to  some  extent  in  the  laboratory.  With  soluble 
sulphides  it  gives  an  intense  violet  color,  but  not  with  hy- 
drogen sulphide.  The  composition  of  the  salt  is  repre- 
sented by  the  formula  Na2Fe(CN)6(NO)  +  2H2O.  The 
free  acid  corresponding  to  this  salt,  and  also  other  salts 
of  the  same  acid,  have  been  made. 

Ferrous  Hydroxide,  Fe(OH)2,  is  formed  when  a  soluble 
hydroxide  is  added  to  a  solution  of  a  ferrous  salt.  It  is 
a  white  precipitate,  but  it  is  usually  obtained  as  a  green- 
ish mass,  as  it  is  very  easily  oxidized  by  the  oxygen  of 


706  INORGANIC  CHEMISTRY. 

the  air  and  that  contained  in  the  solutions.  When  al- 
lowed to  stand  in  contact  with  the  air  it  turns  a  dirty 
green,  and  finally  brown,  being  converted  into  ferric  hy- 
droxide. When  heated  in  the  air  it  loses  water,  and 
takes  up  oxygen,  forming  ferric  oxide. 

Ferrous  Oxide,  FeO,  is  formed  by  passing  hydrogen 
over  ferric  oxide  heated  to  300°.  It  is  a  black  powder, 
which  takes  up  oxygen  from  the  air,  and  is  converted  into 
the  oxide  Fe2O3. 

Ferric  Hydroxide,  Fe(OH)3. — This  compound  is  formed 
most  readily  by  adding  ammonia  to  a  solution  of  a  ferric 
salt,  when  it  appears  as  a  voluminous  brownish-red  pre- 
cipitate. When  filtered,  washed,  and  dried,  its  compo- 
sition is  not  changed.  If  heated  at  100°,  or  if  the  solution 
is  boiled  for  some  time,  it  loses  water,  and  forms  com- 
pounds of  the  formulas  FeO.OH,  Fe2O(OH)4,  etc.  The 
latter  is  derived  from  the  normal  hydroxide  as  repre- 
sented in  the  equation 

2Fe(OH)3  =  Fe20(OH)4  +  H2O. 

The  mineral  pyrosiderite  is  the  hydroxide  FeO.OH. 
Brown  iron  ore  is  Fe4O3(OH)6 ;  and  bog  iron-ore  is 
Fe2O(OH)4.  All  of  these  are  derivatives  of  the  normal 
hydroxide.  The  normal  hydroxide  differs  from  alumin- 
ium hydroxide  in  the  fact  that  it  has  no  acid  properties. 
Therefore,  if  the  two  hydroxides  are  treated  together  with 
a  caustic  alkali  only  the  aluminium  hydroxide  dissolves. 
The  compound  FeO.OH,  corresponding  to  A1O.OH  and 
CrO.OH,  yields  salts  under  some  circumstances.  Thus 

a  calcium  salt,  -p6Q |*Q>Ca,  is  formed  by  heating  together 

ferric  oxide  and  lime  to  a  high  temperature.  In  compo- 
sition this  is  plainly  analogous  to  the  spinels.  Magnetic 
oxide  of  iron  or  magnetite  is  believed  to  be  the  corre- 
sponding ferrous  salt,  -p,6Q *Q>Fe.  Franklinite  also  is  a 
salt  of  the  same  order,  containing  zinc.  It  is  essentially 
a  zinc  salt,  of  the  formula  -p6Q*Q>Zn,  but  some  of  the 
zinc  is  replaced  by  iron  and  manganese. 


OXIDES  OF  IRON.  707 

Ferrous-Ferric  Oxide,  Fe3O4. — As  stated  above,  this 
compound  is  regarded  as  analogous  to  the  spinels,  and 
as  the  ferrous  salt  of  the  acidic  hydroxide  FeO.OH,  as 
represented  in  the  equation  (FeO.O)2Fe.  It  is  found  in 
nature  as  the  mineral  magnetite,  and  loadstone,  which 
occurs  in  Sweden,  Norway,  and  elsewhere.  It  is,  further, 
formed  when  iron  is  burned  in  oxygen,  and  when  water 
is  passed  over  red-hot  iron.  Some  of  the  magnetite 
which  occurs  in  nature  has  the  power  to  attract  iron,  or 
is  magnetic. 

Soluble  Ferric  Hydroxide  is  formed  when  a  solution  of 
ferric  chloride  or  ferric  acetate  is  treated  with  ferric  hy- 
droxide, and  the  solution  thus  formed  dialyzed  (see  page 
421).  The  ferric  salts  pass  through  the  membrane,  and 
the  ferric  hydroxide  remains  in  solution  in  water,  form- 
ing a  deep-red  liquid.  It  is  used  in  medicine.  Small 
quantities  of  salts  cause  the  precipitation  of  ferric  hy- 
droxide from  the  solution. 

Ferric  Oxide,  Fe2O3,  is  found  in  nature,  and  is  known 
as  hematite,  forming  one  of  the  most  valuable  ores  of 
iron.  It  can  be  made  in  the  laboratory  by  igniting  the 
hydroxide.  As  hematite,  it  is  a  black,  crystallized  sub- 
stance with  a  high  lustre.  Otherwise  it  has  a  red  or  a  red- 
dish-brown color.  The  oxide  found  in  nature  and  that 
which  has  been  strongly  ignited  are  very  difficultly  solu- 
ble in  acids.  In  the  preparation  of  fuming  sulphuric 
acid  by  heating  ferrous  sulphate  (see  page  219)  there  is 
left  a  residue  of  ferric  oxide  known  as  rouge,  which  is 
used  as  a  red  pigment  and  as  a  polishing  powder.  A 
specially  fine  variety  of  rouge  for  polishing  is  manufac- 
tured by  heating  ferrous  oxalate,  FeC2O4,  in  contact  with 
the  air. 

Ferrous  Sulphide,  FeS,  is  formed  by  direct  union  of 
iron  and  sulphur  when  the  two  are  heated  together.  It 
is  manufactured  by  heating  iron  filings  and  sulphur  to- 
gether in  a  crucible.  The  pure  compound  is  yellow  and 
crystalline.  When  heated  in  contact  with  the  air  it  is 
oxidized  to  ferrous  sulphate,  if  the  temperature  is  not 
too  high.  At  a  higher  temperature  the  products  are 
sulphur  dioxide  and  ferric  oxide.  When  a  solution  of  a 


708  INORGANIC  CHEMISTRY. 

ferrous  salt  is  treated  with  ammonium  sulphide,  ferrous 
sulphide  is  precipitated  as  a  black  powder.  When  a 
ferric  salt  is  treated  with  ammonium  sulphide  it  is  re- 
duced to  the  ferrous  condition,  and  then  ferrous  sulphide 
is  precipitated : 

Fe2(S04)3  +    (NH4)2S  =  2FeSO4  +(NH4)2SO4  +  S  ; 
2FeS04    +  2(NH4)2S  =  2FeS      +  2(NH4)3SO4. 

The  sulphide  thus  obtained  oxidizes  readily  in  the  air, 
and  forms  the  sulphate.  The  compact  variety  is  used 
in  making  hydrogen  sulphide  (which  see). 

Ferric  Sulphide,  Pe2S3,  is  analogous  to  ferric  oxide, 
Fe2O3.  It  is  formed  artificially  by  heating  iron  and  sul- 
phur together  in  the  proper  proportions. 

Just  as  there  are  salts  derived  from  the  hydroxide 
FeO.OH,  so  there  are  salts  which  are  derived  from  the 
corresponding  sulphide  FeS.SH.  The  potassium,  sodium, 
and  some  other  salts  are  obtained  artificially.  Chalco- 
pyrite  is  apparently  the  cuprous  salt  SFe-S-Cu  or 
FeCuS2. 

Ferrous  Carbonate,  FeCO3. — This  salt  occurs  in  nature 
as  spathic  iron  or  siderite.  It  crystallizes  in  forms 
similar  to  those  of  calc  spar  or  calcium  carbonate  CaCO3. 
Like  this,  further,  it  dissolves  in  water  which  contains 
carbon  dioxide,  and  is  therefore  contained  in  natural 
waters  which  come  in  contact  with  it.  When  a  solution 
of  a  ferrous  salt  is  treated  with  a  soluble  carbonate  a 
white  precipitate  is  formed,  which  is  ferrous  carbonate ; 
but  in  contact  with  the  air  this  is  rapidly  oxidized  and 
decomposed,  leaving  ferric  hydroxide,  which  with  car- 
bonic acid  does  not  form  a  salt.  In  this  respect  ferric 
hydroxide  acts  like  alurninic  and  chromic  hydroxides, 
and  therefore  when  a  soluble  carbonate  is  added  to  a 
solution  of  a  ferric  salt  the  hydroxide  and  not  ferric  car- 
bonate is  thrown  down. 

Ferrous  Sulphate,  FeSO4. — This  important  compound 
is  manufactured  on  the  large  scale  by  the  spontaneous 
oxidation  of  pyrite  in  contact  with  the  air,  and  by  dis- 
solving iron  in  sulphuric  acid.  It  is  frequently  called 


FERROUS  SULPHATE.  709 

"  green  vitriol"  (see  p.  596),  and  more  commonly  "  cop- 
peras." Under  ordinary  conditions  it  crystallizes  in 
transparent,  green,  monoclinic  crystals  with  seven  mole- 
cules of  water,  just  as  zinc  sulphate,  magnesium  sulphate, 
etc.,  do ;  and  when  heated,  six  of  these  are  given  off 
readily,  while  the  last  is  given  off  with  difficulty — a  fact 
which  makes  it  appear  probable  that  the  salt  is  a  deriva- 
tive of  tetrahydroxyl- sulphuric  acid,  as  represented  in 

( (OH), 

the  formula  OS-<  O  -ci  .  While  it  ordinarily  crystal- 
lizes in  monoclinic  crystals,  it  takes  the  rhombic  form 
if  its  supersaturated  solution  is  touched  with  a  crystal 
of  zinc  sulphate.  It  also  crystallizes  in  the  triclinic  sys- 
tem with  five  molecules  of  water,  like  cupric  sulphate,  if 
a  crystal  of  the  latter  salt  is  placed  in  its  concentrated 
solution.  The  salt  undergoes  change  when  exposed  to 
the  air,  being  converted  into  a  compound  containing 
ferric  sulphate,  Fe2(SO4)3,  and  ferric  hydroxide,  or  more 
probably  a  basic  ferric  sulphate,  Fe3(SO4)3(OH)3. 

6FeS04  +  30  +  3H2O  =  2Fe3(SO4)3(OH)3,  or 
6FeSO4  +  3O  +  3H2O  =  2Fe2(SO4)3  +  2Fe(OH)3. 

The  same  change  takes  place  when  a  solution  of  ferrous 
sulphate  is  exposed  to  the  air.  When  treated  with  oxid- 
izing agents  in  the  presence  of  sulphuric  acid  it  is  com- 
pletely converted  into  ferric  sulphate  : 

2FeS04  +  H2S04  +  O  =  Fe,(SO4),  +  H,O. 

Like  other  soluble  ferrous  salts  it  absorbs  nitric  oxide, 
and  when  the  solution  of  the  unstable  compound  is  heated 
the  nitric  oxide  is  given  off.  Ferrous  sulphate  is  used  in 
dyeing,  in  the  manufacture  of  ink,  etc.;  and  as  a  deodor- 
izer. With  sulphates  of  the  alkalies  ferrous  sulphate 
forms  double  salts,  such  as  FeK2(SO4)2  +  6H2O,  Fe(NH4), 
(SO4)2  -f-  6H2O,  etc.  These  are  not  as  easily  oxidized  as 
the  simple  salt,  and  are  convenient  in  the  laboratory 
when  a  pure  ferrous  salt  is  wanted.  It  is  a  fact  worthy 
of  special  notice,  that  while  ferrous  sulphate  crystallizes 


710  INORGANIC  CHEMISTRY. 

with  seven  molecules  of  water,  these  salts  contain  only 
six  molecules,  and  all  of  this  is  easily  given  off  when 
the  salts  are  heated.  It  appears,  therefore,  that  these 
double  salts  are  formed  from  ferrous  sulphate  by  re- 
placing one  molecule  of  the  water  by  a  molecule  of  some 
sulphate.  This  is  clear  if  ferrous  sulphate  and  the  other 
salts  are  regarded  as  salts  of  tetrahydroxyl-sulphuric 
acid.  We  should  then  have  the  relation  between  the 
double  sulphate  and  the  simple  ones  as  represented  in 
the  formulas  below : 


H         HO 
H         HO 


O 


OK 
OK 


Ferrous  sulphate  Potassium  sulphate 

Fe<°>S<g>S<gf  =  FeK,(S04), 

Ferrous-potassium  sulphate 

Ferric  Sulphate,  Fe2(SO4)3. — This  salt,  as  stated  in  the 
last  paragraph,  is  formed  by  oxidation  of  ferrous  sul- 
phate. It  is  also  formed  by  dissolving  ferric  oxide  or 
hydroxide  in  sulphuric  acid.  When  the  solution  is 
evaporated  the  salt  remains  behind  as  a  white,  anhydrous 
mass.  It  readily  forms  basic  salts,  the  composition  of 
which  is  not  positively  known.  With  the  sulphates  of 
the  alkali  metals  it  forms  double  salts,  which  are  per- 
fectly analogous  to  alum,  and  are  known  as  the  iron  alums  ; 
as,  for  example,  FeK(SO4)2  +  12H.2O,  Fe(NH4)(SO4)2  + 
12H2O,  etc. 

Ferrous  Phosphate,  Fe3(PO4)2,  occurs  in  nature  crystal- 
lized with  eight  molecules  of  water  as  the  mineral  vivi- 
anite.  Both  this  salt  and  ferric  phosphate,  FePO4,  are 
insoluble  and  are  formed  when  solutions  of  ferrous  and 
ferric  salts  are  treated  with  sodium  phosphate. 

Ferric  Acid,  H2FeO4,  is  analogous  in  composition  to 
chromic  and  manganic  acids.  The  acid  itself  is  not  known, 
but  its  potassium  salt  is  formed  when  iron  or  ferric  oxide 
is  heated  with  saltpeter,  or  when  chlorine  is  passed  into 


REACTIONS  OF  IRON  COMPOUNDS.  711 

caustic  potash  containing  ferric  hydroxide  in  suspension : 

Fe(OH)3  +  2KOH  +  3C1  =  K2FeO4  +  H2O  +  3HC1. 

This,  as  well  as  the  other  ferrates,  is  unstable,  the  iron 
tending  to  pass  back  into  the  condition  of  a  ferric  com- 
pound. 

Iron  Bisulphide,  PeS2,  is  not  analogous  to  any  oxygen 
compound  of  iron.  In  it  the  metal  appears  to  be  quad- 
rivalent. The  disulphide  occurs  very  widely  distributed 
and  in  large  quantities  in  nature  as  the  mineral  iron 
pyrites  or  pyrite,  which  crystallizes  in  the  regular  sys- 
tem, and  as  marcasite,  which  crystallizes  in  the  rhombic 
system.  It  can  be  made  artificially,  and  if  crystallized 
it  appears  in  the  form  of  pyrite.  Its  conduct  under  the 
influence  of  heat  has  been  repeatedly  referred  to  in  con- 
nection with  the  roasting  of  iron  and  other  ores.  As 
pyrite  it  has  a  golden-yellow  color,  and  it  has  frequently 
been  taken  for  the  precious  metal  by  those  not  familiar 
with  it.  The  name  "  fool's  gold,"  by  which  it  is  some- 
times popularly  known,  suggests  this  fact. 

Arsenopyrite,  FeAsS,  occurs  in  nature,  and  is  a  valu- 
able source  of  the  element  arsenic ;  for,  as  has  been 
stated  (see  p.  305),  when  it  is  heated  it  gives  off  arsenic, 
and  ferrous  sulphide  is  left  behind. 

Iron  Carbonyls. — When  finely-divided  iron  is  alloweS 
to  cool  in  hydrogen  gas  to  80°,  and  then  treated  with  car- 
bon monoxide,  the  issuing  gas  imparts  a  yellow  color  to 
the  flame  of  a  Bunsen  burner,  and  by  passing  it  through 
a  heated  glass  tube  a  metallic  mirror  is  formed  at  between 
200°  and  350°.  This  mirror  consists  of  iron.  These 
phenomena  are  due  to  the  formation  of  volatile  com- 
pounds of  iron  with  carbon  monoxide,  the  principal  one 
having  the  composition  Fe(CO)6.  (See  Nickel  Carbonyl.) 

Reactions  which  are  of  Special  Value  in  Chemical 
Analysis. — Ferrous  Compounds. —  The  reactions  of  ferrous 
compounds  with  the  soluble  hydroxides  and  carbonates, 
ammonium  sulphide,  potassium  ferricyanide,  and  with  ox- 
idizing  agents  have  been  explained  above.  With  ammo- 
nium salts  ferrous  chloride  forms  double  salts,  which  are 
soluble ;  therefore,  if  ammonium  chloride  is  added  to  a 


712  INORGANIC  CHEMISTRY. 

solution  of  the  salt  ammonia  does  not  precipitate  the 
hydroxide.  Further,  ammonia  does  not  completely  pre- 
cipitate the  hydroxide  from  a  solution  of  a  ferrous  salt,  as 
an  ammonium  salt  is  formed.  By  standing  in  the  air, 
however,  these  solutions  containing  the  double  salts  are 
oxidized,  and  ferric  hydroxide  is  precipitated.  The  re- 
actions with  potassium  cyanide  will  be  understood  from 
what  has  been  said  concerning  the  compounds  of  ferro- 
hydrocyanic  and  ferrihydrocyanic  acids. 

Ferric  Compounds. — The  reactions  of  ferric  compounds 
with  the  soluble  hydroxides  and  carbonates,  ammonium  sul- 
phide, potassium  ferrocyanide,  and  potassium  ferricyanide 
have  been  explained  above.  When  hydrogen  sulphide  is 
passed  through  a  solution  of  a  ferric  salt,  reduction  to 
the  corresponding  ferrous  salt  takes  place,  and  sulphur 
separates,  which  gives  the  solution  a  milky  appearance : 

Fe2(S04)3  +  H2S  =  2FeS04  +  H2SO4  +  S  ; 
2FeCl3      +  H2S  =  2FeCl2  +  2HC1  +  S. 

When  a  neutral  solution  of  a  ferric  salt  is  treated  with 
suspended  barium  carbonate  the  iron  is  precipitated  as  the 
hydroxide. 

When  a  neutral  solution  of  a  ferric  salt  is  treated  with 
acetate  of  potassium  or  sodium  it  turns  dark  red,  in  conse- 
quence of  the  formation  of  ferric  acetate  which  remains 
in  solution.  When  the  solution  is  boiled  the  acetate 
breaks  down  into  acetic  acid  and  ferric  hydroxide,  which 
is  precipitated : 

FeCl3  +  3NaC2H302  =  Fe(C2H8O2)3  +  3NaCl ; 
Fe(C2H302)3  +  3H20  =  Fe(OH)3        +  3C2H4Oa. 

When  potassium  sulphocy.anate,  KCNS,  is  added  to  a  solu- 
tion of  a  ferric  salt  a  blood-red  color  is  produced.  This 
occurs  even  in  extremely  dilute  solutions  of  ferric  salts. 

The  borax  bead  is  colored  bottle-green  in  the  reducing 
flame,  and  brown-red  to  yellowish  red  in  the  oxidizing 
flame,  when  treated  with  compounds  of  iron. 

COBALT,  Co  (At.  Wt.  58.49). 


COBALT.  713 

General. — As  stated  in  the  remarks  introductory  to  the 
iron  group,  cobalt,  like  nickel,  preferably  forms  com- 
pounds which  are  analogous  to  ferrous  compounds.  Itr 
however,  forms  a  few  which  are  analogous  to  ferric  com- 
pounds, its  power  in  this  direction  being  greater  than 
that  of  nickel.  Its  salts  form  a  great  variety  of  com- 
pounds with  ammonia,  and  these  have  been  extensively 
studied. 

Occurrence  and  Preparation. — Cobalt  occurs  in  nature, 
almost  always  in  company  with  nickel.  The  principal 
minerals  containing  it  are  smaltite,  CoAs2,  and  cobaltite, 
CoS2.CoAs2.  In  each  of  these  a  part  of  the  cobalt  is  re- 
placed by  iron,  and,  generally,  some  nickel.  By  roasting 
and  melting  the  ores  in  blast-furnaces  they  are  partly 
purified.  The  product  is  dissolved  in  hydrochloric  acid, 
and  treated  with  a  small  quantity  of  calcium  hypochlorite 
and  hydroxide  for  the  purpose  of  removing  iron  and 
arsenic  ;  then  with  hydrogen  sulphide  to  remove  copper 
and  bismuth ;  and  finally  with  calcium  hypochlorite 
when  cobaltic  hydroxide,  Co(OH)3,  is  precipitated.  This 
is  readily  converted  into  the  oxide,  Co2O3,  from  which 
the  metal  can  be  prepared  by  heating  it  in  a  current  of 
hydrogen.  It  is  also  obtained  by  heating  the  oxalate  to 
a  sufficiently  high  temperature. 

Properties. — Cobalt  has  a  silver-white  color,  with  a 
slight  cast  of  red.  It  is  harder  than  iron,  and  melts  at  a 
somewhat  lower  temperature  ;  is  tenacious  ;  and  has  the 
specific  gravity  8.9.  It  dissolves  in  nitric  acid. 

Cobaltous  Chloride,  CoCl2,  is  formed  by  heating  cobalt 
in  chlorine  gas,  and  in  solution  by  treating  cobalt 
carbonate  with  hydrochloric  acid.  From  the  solution 
it  crystallizes  in  dark-red  prisms  of  the  composition 
CoCL,  +  6H2O.  The  anhydrous  salt  is  blue.  When  the 
blue  salt  is  treated  with  water  it  turns  red,  and  when  the 
red  salt  is  heated  it  turns  blue.  This  difference  in  color 
between  the  anhydrous  and  the  hydrated  salts  is  charac- 
teristic of  cobalt  salts.  If  marks  are  made  on  paper 
with  a  dilute  solution  of  one  of  the  salts  the  color  is  not 
perceptible.  If,  however,  the  paper  is  held  before  a 
fire,  the  salt  loses  water  and  turns  blue,  and  as  the  blue 


714  INORGANIC  CHEMISTRY. 

is  more  intense  than  the  red,  it  is  visible.  When  the  salt 
becomes  moist  again  it  becomes  invisible.  This  is  the 
basis  for  the  preparation  of  the  so-called  sympathetic  inks. 

Cobaltous  Hydroxide,  Co(OH)2,  is  formed  as  a  red  pre- 
cipitate when  a  soluble  hydroxide  is  added  to  a  cobalt- 
ous  salt,  and  the  blue  precipitate,  which  is  first  formed 
and  which  is  a  basic  salt,  is  allowed  to  stand.  It  is  oxid- 
ized by  contact  with  the  air,  forming  cobaltic  hydroxide^ 
which  breaks  down  into  cobaltic  oxide,  Co2O3. 

Cobaltous  Oxide,  CoO,  is  formed  when  the  correspond- 
ing hydroxide  is  carefully  heated  without  access  of  air. 
"When  heated  in  the  air  it  is  converted  into  cobaltous- 
cobaltic  oxide,  Co3O4,  which  is  analogous  to  ferrous-ferric 
oxide.  This  is  also  formed  when  cobaltic  oxide,  Co2O3, 
is  heated  in  the  air. 

Cobaltic  Hydroxide,  Co(OH)3,  is  formed  when  calcium 
hypochlorite  is  added  to  a  solution  of  a  cobaltous  salt, 
and  is  a  black  powder.  When  heated  it  is  converted 
into  black  cobaltic  oxide,  Co2O3. 

Cobalt  Sulphide,  CoS,  is  the  black  precipitate  which  is 
formed  by  adding  ammonium  sulphide  to  a  cobaltous 
salt.  It  is  not  soluble  in  dilute  acids,  and  differs  from 
ferrous  sulphide  in  this  respect.  Other  sulphides  of 
cobalt  are  those  of  the  formulas  Co3S4  and  CoS2.  The 
former  is  found  in  nature,  and  is  known  as  linnseite. 
The  latter  occurs  in  combination  with  other  sulphides, 
as  in  cobaltite,  CoS2.CoAs2. 

Cyanides. — Cobaltous  cyanide  is  an  insoluble  dirty-red 
compound  which  is  formed  when  potassium  cyanide  is 
added  to  a  solution  of  a  cobalt  salt.  It  dissolves  in  an 
excess  of  potassium  cyanide,  forming  a  double  cyanide, 
K4Co(CN)6,  which  is  analogous  to  potassium  ferrocya- 
nide.  When  this  solution  is  boiled  the  cyanide  is  oxi- 
dized, forming  a  compound  analogous  to  potassium 
ferricyanide,  thus : 

2K4Oo(ON).  +  H2O  +  O  =  2K3Co(CN)6  +  2KOH. 

This  acts  like  the  corresponding  iron  compound^  The 
cobalt  is  not  precipitated  from  it  by  ammonium  sulphide 


COMPOUNDS  OF  AMMONIA  WITH  SALTS  OF  COBALT.  715 

or  sodium  hydroxide.  This  conduct  towards  potassium 
cyanide  distinguishes  cobalt  from  nickel  salts. 

Smalt. — The  beautiful  pigment  known  by  this  name  is 
essentially  a  cobalt  glass  in  which  cobalt  takes  the  place 
of  calcium.  It  is  made  by  heating  compounds  of  cobalt 
with  quartz  and  potassium  carbonate.  The  glass  thus 
formed  is  powdered  very  finely  and  used  as  a  pigment. 
It  does  not  change  color  in  the  sunlight,  and  is  not  af- 
fected by  acids  nor  by  alkalies. 

Compounds  of  Ammonia  with  Salts  of  Cobalt. — In  gen- 
eral, when  solutions  of  cobalt  salts  in  ammonia  are  ex- 
posed to  the  air  they  undergo  oxidation,  and  complicated 
salts  are  formed.  In  the  case  of  cobaltous  chloride, 
the  first  product  formed  is  one  of  the  composition 
Co(NH3)3Cls  +  HaO,  which  is  known  as  dichro-cobaltic 
chloride.  At  the  same  time  another  compound  of  the 
composition  Co(NH3)4Cl3  -|-  H2O,  known  as  praseo-cobaltic 
chloride,  is  formed.  If  a  solution  of  cobaltous  chloride  in 
concentrated  ammonia  is  allowed  to  stand  longer  than  is 
required  to  form  the  preceding  compound,  or  if  an  oxid- 
izing agent  is  used,  the  product  has  the  composition 
Co(NH3)6Cl3,  and  is  ^nown  as  purpureo-cobaltic  chloride. 
And,  finally,  by  further  action  of  oxidizing  agents,  luteo- 
cobaltic  chloride,  Co(NH3)6Cl3,  is  formed.  It  will  be  ob- 
served that  these  compounds  form  a  series,  the  members 
of  which  differ  from  one  another  by  NH3  or  a  multiple : 

Co(NH,)sCl,;    Co(NH,)4Cl,;    Co(NH.)6Cl, ;    Co(NH,).Cl, 

They  may  be  regarded  as  made  up  of  cobaltic  chloride 
and  different  numbers  of  molecules  of  ammonia.  In 
regard  to  the  first  one,  it  is  simplest  to  consider  it  as 
analogous  to  the  mercur-ammonium  compounds  (see 
page  633).  Accordingly,  it  is  usually  represented  as 
derived  from  three  molecules  of  ammonium  chloride  by 
the  substitution  of  a  trivalent  atom  of  cobalt  for  three 
hydrogen  atoms. 

When  a  solution  of  cobaltous  chloride  in  ammonia, 
which  has  become  red  by  contact  with  the  air,  is  treated 
at  the  ordinary  temperature  with  concentrated  hydro- 
chloric acid,  a  brick-red  precipitate  of  roseo-cobaltic  chlo- 


716  INORGANIC  CHEMISTRY. 

ride,  which  has  the  same  composition  as  purpureo-co- 
baltic  chloride,  Co(NH3)BCl3,  with  a  molecule  of  water, 
is  formed. 

If  in  the  above  reactions  the  nitrate  or  sulphate  of 
cobalt  is  used,  nitrates  and  sulphates  corresponding  to 
the  chlorides  mentioned  are  obtained. 

Thus : 

Hoseo-salts.  Luteo-salts. 

Co(NH3)6013  Co(NH3)6Cl3 

Co(NH3)6(N03)3  Co(NH3)6(N03)3 

[Co(NH3)Ja(S04)3  [Co(NH3)J3(S04)3 

NICKEL,  Ni  (At.  Wt.  58.24). 

General. — Nickel  differs  from  cobalt  in  respect  to  the 
difficulty  with  which  it  forms  nickelic  compounds  or 
those  in  which  it  is  trivalent.  At  the  same  time  it  does 
form  an  oxide  of  the  composition  Ni2O3,  and  the  corre- 
sponding hydroxide,  Ni(OH)3.  In  all  other  compounds 
it  is  bivalent,  the  compounds  being  analogous  to  ferrous 
compounds. 

Occurrence  and  Preparation. — Nickel  occurs  native  in 
meteorites.  The  principal  minerals  containing  it  are 
the  arsenide,  NiAs,  known  as  niccolite,  and  the  sulph- 
arsenide,  NiSAs  or  NiS2.NiAs2,  known  as  gersdorffite. 
From  the  ores  the  oxide  is  obtained  in  the  same  way 
that  cobalt  oxide  is  obtained  from  its  ores.  This  is  then 
pressed  in  the  form  of  small  cubes,  mixed  with  charcoal 
powder,  and  ignited.  The  commercial  metal  is  generally 
found  in  the  form  of  these  cubes. 

Properties. — Nickel  is  a  white  metal  with  a  slight  cast 
of  yellow.  It  is  very  hard,  and  capable  of  a  high  polish. 
The  metal  in  its  ordinary  condition  is  brittle,  but  after 
treatment  with  a  little  magnesium  it  becomes  very  mal- 
leable. Its  specific  gravity  is  8.9,  and  it  melts  at  a  high 
temperature.  It  is  not  changed  in  the  air ;  it  dissolves 
slowly  in  hydrochloric  and  sulphuric  acids,  and  readily 
in  nitric  acid.  Like  iron,  it  is  magnetic. 

Alloys. — Alloys  of  nickel  are  extensively  used.  A.r- 
gentan  or  German  silver  consists  of  copper,  zinc,  and  nickel. 
Various  nickel  alloys  are  used  for  making  coins.  The 
5  and  3  cent  pieces  in  the  United  States  are  made  of  an 


COMPOUNDS  OF  NICKEL.  717 

alloy  consisting  of  25  per  cent  nickel  and  75  per  cent 
copper.  In  Switzerland,  and  Belgium  also,  nickel  coins 
are  used. 

Other  Applications  of  Nickel. — Besides  as  a  constituent 
of  important  alloys,  nickel  is  extensively  used  at  present 
in  nickel-plating.  Iron  objects  are  covered  with  a  thin 
layer  of  the  metal  for  the  purpose  of  protecting  them 
from  rusting.  The  plating  is  accomplished  as  silver- 
plating  and  copper-plating  are — by  means  of  electrolysis, 
-a  bath  of  nickel-ammonium  sulphate  being  used. 

Nickelous  Chloride,  NiCl2,  crystallizes  from  aqueous 
solution  with  six  molecules  of  water,  and  the  crystals  are 
green.  When  the  water  of  crystallization  is  driven  off 
they  become  yellow.  In  general,  nickel  salts,  with  their 
water  of  crystallization,  are  green,  and  in  the  anhydrous 
condition  they  are  yellow. 

Nickelous  Hydroxide,  Ni(OH)2,  is  formed  when  a  nickel 
salt  is  treated  with  a  soluble  hydroxide,  and  is  a  green 
insoluble  substance.  When  heated  it  is  converted  into 
the  green  oxide,  NiO. 

Nickelic  Hydroxide,  Ni(OH)3,  is  precipitated  as  a  black 
powder  when  a  solution  of  a  nickel  salt  is  treated  with 
sodium  hypochlorite. 

Cyanides. — When  potassium  cyanide  is  added  to  a  so- 
lution of  a  nickel  salt,  nickel  cyanide,  Ni(CN)a,  is  precipi- 
tated as  a  greenish-white  substance.  With  an  excess  of 
potassium  cyanide  this  forms  the  salt,  Ni(CN)2.2KCN, 
which,  owing  to  the  fact  that  nickelous  salts  are  not 
•converted  into  nickelic  salts  by  oxidation,  does  not 
undergo  change  when  boiled  with  potassium  cyanide. 
When  hydrochloric  acid  is  added  to  a  solution  of  the 
double  cyanide,  nickelous  cyanide  is  precipitated.  If 
boiled  with  precipitated  mercuric  oxide  the  double  cy- 
anide is  decomposed  and  nickel  oxide  is  thrown  down. 

Nickel  Carbonyl,  Ni(CO)4. — This  interesting  compound 
is  formed  when  finely-divided  nickel,  such  as  is  obtained 
by  reducing  nickel  oxide  by  hydrogen  at  about  400°,  is 
allowed  to  cool  in  a  slow  current  of  carbon  monoxide. 
A  gas  is  formed  which  can  easily  be  condensed,  its 
boiling-point  being  43°.  At  —25°  it  solidifies  forming 


718  INORGANIC  CHEMISTRY. 

needle-shaped  crystals.  When  the  gas  is  passed  through 
a  heated  tube  pure  nickel  is  deposited.  Advantage  of 
this  fact  is  taken  for  the  purpose  of  preparing  pure 
nickel  on  the  large  scale.  Cobalt  does  not  form  a  com- 
pound of  this  kind. 

Reactions  of  Cobalt  and  Nickel  which  are  of  Special 
Value  in  Chemical  Analysis. — The  reactions  with  the  sol- 
uble hydroxides  have  been  explained.  With  ammonium 
sulphide  both  give  black  sulphides,  which  are  not  easily 
dissolved  by  dilute  hydrochloric  acid.  From  solutions 
of  the  acetates  hydrogen  sulphide  precipitates  the  sul- 
phides. Nickel  sulphide  is  slightly  soluble  in  ammonium 
sulphide,  and  the  solution  has  a  brownish-yellow  color. 

The  action  of  the  hypochlorites  upon  solutions  of 
nickel  and  cobalt  salts  has  been  explained  above.  The 
reactions  with  potassium  cyanide  have  also  been  ex- 
plained. These  furnish  a  good  method  for  separating, 
the  two  metals. 

When  a  solution  of  potassium  nitrite  is  added  to  a  solu- 
tion of  a  cobalt  salt  containing  free  acetic  acid  or  nitria 
acid,  a  precipitate  of  cobaltic  potassium  nitrite  is  formed. 
This  is  a  compound  of  cobaltic  nitrite,  Co(NO2)3,  and  potas- 
sium nitrite,  of  the  composition  Co(NO2)3.3KNO2.  The 
formation  involves  oxidation  of  the  cobaltous  salt,  and 
this  is  effected  by  some  of  the  nitrogen  trioxide  which  is 
set  free.  Thus  with  the  chloride  the  action  may  be  rep- 
resented as  follows : 

CoCl2  +  7KNO2  +  2C2H402  = 

2KC1  +  Co(N02)3.3KN02  +  2KC2H3O2  +  H2O  +  NO. 

Nickel  does  not  form  a  similar  compound  of  a  nickelic 
salt,  but  simply  forms  a  double  nitrite,  containing  the 
nickelous  salt  Ni(NO2)2.4KNO2. 

Cobalt  compounds  color  the  bead  of  microcosmic  salt 
blue  both  in  the  reducing  and  oxidizing  flame.  Nickel 
colors  it  reddish  brown  in  the  oxidizing  flame  when  hot, 
and  pale  yellow  when  cold.  In  the  reducing  flame  it  ia 
gray. 


CHAPTER  XXXIV. 

ELEMENTS  OF  FAMILY  VIII,  SUB  GROUP  B  j 
RUTHENIUM— RHODIUM— PALLADIUM. 

ELEMENTS  OF  FAMILY  VIII,  SUBGROUP  C 
OSMIUM— IRIDIUM— PLATINUM. 


General. — Comparing  the  members  of  the  three  sub- 
groups of  Family  VIII,  with  reference  to  their  atomic 
weights  and  specific  gravities,  we  have  the  following  re- 
in a  rkable  table: 


Fe 

Co 

Ni 

At.  Wt.    55.6 

At.  Wt.    58.49 

At.  Wt.    58.24 

Sp.  Gr.      7.8 

Sp.  Gr.      8.5 

Sp.  Gr.      8.8 

Ru 

Rh 

Pd 

At.  Wt.  100.91 

At.  Wt.  102.23 

At.  Wt.  105.56 

Sp.  Gr.    12.26 

Sp.  Gr.    12.1 

Sp.  Gr.    11.5 

Os 

Ir 

Pt 

At.  Wt.  189.55 

At.  Wt.  191.66 

At.  Wt.  193.41 

Sp.  Gr.    22.48 

Sp.  Gr.    22.42 

Sp.  Gr.    21.50 

It  will  be  observed  that  the  atomic  weights  and  specific 
gravities  of  the  members  of  each  sub-group  are  approxi- 
mately the  same.  But  just  as  there  is  a  gradual  change 
in  the  chemical  conduct  as  we  pass  from  iron  to  nickel 
in  the  iron  group,  so  a  similar  gradation  of  properties  is 
observed  in  the  other  two  groups.  As  far  as  the  variety 
of  compounds  which  the}'  form  is  concerned,  ruthenium 
and  osmium  are  more  like  iron  than  they  are  like  rho- 
dium and  iridium.  Further,  rhodium  and  iridium  re- 
semble each  other,  as  regards  the  variety  of  their  com- 
pounds, more  closely  than  they  resemble  palladium  and 
platinum,  and  a  similar  resemblance  is  noticed  between 
palladium  and  platinum.  These  relations  will  appear 

(719) 


720  INORGANIC  CHEMISTRY. 

more  clearly  if  the  formulas  of  some  of  the  principal 
compounds  of  the  elements  under  consideration  are 
placed  together  in  a  table. 

Ku  and  Os  Bh  and  Ir                     Pd  and  Pt 

Eu04        OsO4  EhO2      IrOa            PdOa     PtO, 

HKuO4  KhaO3     Ir2O8           PdO       PtO 

H2Ku04    HaOsO4  EhO       IrO             Pd3O 

EuOa        OsO2  IrCl4           PdCl4     PtCl4 

Ku2O8       Os2O3  BhCls    IrCl3           PdCl2     PtCl, 

EuO         OsO  IrCla 

EuCl4       OsCl4 

EuCl3       OsCl3 

EuCla       OsCla 

The  elements  ruthenium  and  osmium  have  a  more  acidic 
character  than  the  others  ;  just  as  iron  has  a  more  acidic 
character  than  cobalt  and  nickel.  Euthenium  forms  not 
only  ruthenious  acid,  H2EuO4,  which  is  analogous  to 
ferric,  chromic,  and  manganic  acids,  but  also  perru- 
thenious  acid  analogous  to  permanganic  acid.  Osmium 
forms  osmious  acid,  H2OsO4,  but  apparently  no  peros- 
mious  acid.  The  highest  known  oxides  are  derived  from 
these  elements.  These  are  the  tetroxides,  EuO4  and 
OsO4,  in  which  the  elements  appear  to  be  octovalent. 
Neither  rhodium  nor  iridium  forms  acids.  As  far  as  their 
oxides  and  chlorides  are  concerned,  they  suggest  manga- 
nese more  than  any  other  element,  but  their  oxides  have 
only  weak  basic  properties.  Passing  finally  to  the  last 
pair,  palladium  and  platinum,  we  find  that  they  have  not 
the  power  to  form  oxides  of  the  general  formula  M2O3> 
but  that  they  act  as  the  members  of  Family  IV  do — either 
as  bivalent  or  quadrivalent  elements.  Palladium,  to  be 
sure,  forms  a  compound,  the  sub-oxide  Pd2O,  which  is 
like  the  oxide  of  silver,  Ag2O,  and  in  which  it  appears  to 
be  univalent.  In  fact  the  members  of  Family  VIII  form 
the  connecting  link  between  the  members  of  Family  VII 
and  those  of  Family  I.  In  manganese,  as  we  have  seen, 
a  maximum  of  power  is  reached  as  far  as  the  valence  is 
concerned.  It  forms  compounds  in  which  it  appears  to 


GENERAL  IN  REGARD  TO  THE  PLATINUM  METALS.     721 

be  septivalent,  sexivalent,  quadrivalent,  trivalent,  and 
bivalent.  When  we  pass  to  iron,  however,  we  find  that 
it  is  not  septivalent.  In  its  most  complex  compounds 
this  element  is  sexivalent,  as  in  ferric  acid,  H2FeO4, 
but  it  acts  preferably  as  a  trivalent  or  a  bivalent 
element.  Then,  further,  as  we  have  seen,  cobalt  forms 
a  few  compounds  in  which  it  is  trivalent,  but  it  is 
generally  bivalent,  and  nickel  is  scarcely  ever  trivalent. 
In  its  compounds  nickel  resembles  copper  in  the  cupric 
compounds,  and  copper  is  the  next  element  in  the  order 
of  increasing  atomic  weights.  'But  copper  has  an  ad- 
ditional power  which  allies  it  to  the  members  of  Group 
A,  Family  I.  It  acts  as  a  univalent  element  in  the 
cuprous  compounds. 

Now,  in  the  same  way,  there  is  an  increase  in  the 
complexity  of  the  compounds  formed  by  the  elements,  as 
we  pass  from  zirconium,  to  niobium,  to  molybdenum,  and 
below  manganese  in  Family  VII  we  should  expect  to  find 
an  element  forming  compounds  which  in  general  resem- 
ble those  of  manganese,  and  leading  up  to  the  octovalent 
element  ruthenium.  Considering  the  relations  between 
iron  and  ruthenium,  one  is  tempted  to  suspect  that 
this  unknown  element  may  exhibit  a  valence  of  nine 
in  some  unstable  compounds.  While  ruthenium  is  oc- 
tovalent in  its  highest  oxide,  it  is  also  septivalent  in 
HEuO4,  sexivalent  in  H2KuO4,  quadrivalent  in  EuO2, 
trivalent  in  Ru2O3,  and  bivalent  in  EuO.  Rhodium, 
however,  is  only  quadrivalent,  trivalent,  and  bivalent ; 
and  palladium  is  quadrivalent,  bivalent,  and  univalent. 
Just  as  nickel  leads  naturally  to  copper,  so  palladium 
leads  naturally  to  silver. 

In  regard  to  the  series  to  which  osmium,  iridium,  and 
platinum  belong,  not  as  much  is  known  as  in  regard  to 
the  series  just  referred  to,  though  the  three  elements 
themselves  have  been  carefully  studied.  There  is  here 
observed  the  same  falling  off  of  valence  power  from  os- 
mium to  platinum  ;  and* just  as  nickel  leads  to  copper, 
and  palladium  leads  to  silver,  so  platinum  leads  natu- 
rally to  gold  in  Family  I. 


INORGANIC  CHEMISTRY. 


THE  PLATINUM  METALS. 

The  six  elements  of  Sub-Groups  B  and  C,  Family 
VIII,  are  generally  grouped  together  and  spoken  of  as 
the  platinum  metals.  They  occur  together  in  nature, 
and  almost  always  in  alloys,  into  the  composition  of 
which  all  enter.  The  chief  constituent  is  platinum, 
which  is  present  to  the  extent  of  50  to  80  per  cent,  and 
over.  The  alloys  occur  in  only  a  few  localities,  in  the 
Ural  Mountains,  in  California,  Australia,  Borneo,  and  a 
few  other  places,  and  form  small  pieces  which  are  mixed 
with  sand  and  earth.  They  generally  contain  also  gold, 
iron,  and  copper.  Palladium  occurs,  further,  in  a  gold 
ore  which  is  found  in  Brazil. 

Metallurgy. — The  process  for  obtaining  the  metals 
from  the  ores  is  based  mainly  upon  the  following  facts : 
(1)  Gold  is  soluble  in  dilute  aqua  regia,  while  platinum 
requires  concentrated  aqua  regia;  (2)  platinic  chloride, 
PtCl4,  and  iridium  chloride,  IrCl4,  form,  with  ammonium 
chloride,  difficultly  soluble  compounds  of  the  formulas 
(NH4)2PtCl6(PtCl4.2NH4Cl)  and  (NH4)  JrCl6(IrCl4.2NH4Cl). 
When  these  compounds  are  ignited,  they  are  completely 
decomposed,  and  the  metals  are  left  behind.  When, 
therefore,  platinum-ore  has  been  freed  as  far  as  possible 
from  sand  and  earth,  it  is  first  treated  with  dilute  aqua 
regia,  which  removes  the  gold,  and  then  with  concen- 
trated aqua  regia,  which  dissolves  the  platinum  together 
with  a  little  iridium,  leaving  an  alloy  of  iridium  and  os- 
mium. 

When  the  solution  thus  obtained  is  treated  with 
ammonium  chloride,  both  metals  are  precipitated ;  and 
when  the  precipitate  is  ignited,  both  metals  are  left  be- 
hind in  the  form  of  a  spongy  mass.  This  consists,  how- 
ever, almost  wholly  of  platinum,  the  amount  of  iridium 
being  very  small. 

KUTHENIUM,  Eu  (At.  Wt.  100.91). 

Preparation. — Euthenium  is  obtained  from  the  residue 
which  is  left  undissolved  when  platinum-ore  is  treated 
with  concentrated  nitro- hydrochloric  acid. 


RUTHENIUM— OSMIUM.  723 

Properties. — When  heated  in  oxygen  it  burns  and  forms 
the  oxide,  EuO2.  It  is  insoluble  in  the  strong  acids,  and 
even  in  nitro-hydrochloric  acid  it  is  almost  insoluble. 
Owing  to  its  power  to  form  salts  of  ruthenious  acid,  it  is 
dissolved  when  heated  with  potassium  hydroxide  and  an 
oxidizing  agent,  such  as  saltpeter  or  potassium  chlorate, 
and  afterwards  treated  with  water. 

Chlorides. — When  heated  in  chlorine  it  forms  the  di- 
chloride,  EuCl2,  and  some  of  the  trichloride,  EuCl3.  The 
tetrachloride,  EuCl4,  is  known  in  combination  with  chlo- 
rides of  the  alkali  metals. 

Oxides. — When  ruthenium  is  heated  with  potassium 
hydroxide  and  saltpeter,  potassium  ruthenite,  K2EuO4,  is 
formed.  The  acid  from  which  this  salt  is  derived  is 
plainly  ruthenious  acid,  H2EuO4,  and  this  is  related  to 
the  oxide,  EuO3.  Neither  the  acid  nor  the  anhydride  is 
known,  however.  When  the  solution  is  treated  with 
chlorine  the  first  product  is  potassium  perruthenite, 
KEuO4,  which  forms  a  dark  green  solution,  and  is  iso- 
morphous  with  potassium  permanganate  and  potassium 
perchlorate.  By  further  treatment  of  the  solution  with 
a  rapid  current  of  chlorine,  ruthenium  peroxide,  EuO4,  is 
formed.  This  is  a  volatile  crystalline  solid,  which  ap- 
parently is  not  acidic.  It  is  easily  reduced  to  the  ses- 
quioxide,  Eu2O3 ;  and,  if  heated,  it  is  decomposed  with 
explosion. 

The  oxides  of  the  formulas  EuO2,  Eu2O3,  and  EuO  are 
not  basic,  and  do  not  dissolve  in  acids. 

OSMIUM,  Os  (At.  Wt.  189.55). 

Preparation. — As  stated  above,  this  element  is  left  un- 
dissolved  in  the  form  of  an  alloy  with  iridium  when  plati- 
num-ore is  treated  with  concentrated  nitro-hydrochloric 
acid.  In  order  to  separate  it  from  the  iridium,  advan- 
tage is  taken  of  the  fact  that  it  forms  a  volatile  peroxide, 
OsO4,  similar  to  that  formed  by  ruthenium,  while  iridium 
does  not. 

Properties. — The  metal  does  not  melt  at  the  highest 
temperatures  reached  artificially.  It  has  the  highest 
specific  gravity  of  all  known  substances  ;  is  easily  oxid- 


724  INORGANIC  CHEMISTRY. 

ized  when  in  finely  divided  condition  ;  and  is  converted 
either  by  the  oxygen  of  the  air  or  by  nitric  acid  into 
osmium  peroxide,  OsO4. 

Chlorides. — The  dicliloride,  OsCl2,  and  the  tetraMoride, 
OsCl4,  are  formed  by  treating  the  metal  with  chlorine* 
The  trichloride,  OsCl3,  is  not  known  in  free  condition. 

Oxides. — The  metal  as  well  as  the  oxides  forms  the  per- 
oxide, OsO4,  when  heated  in  the  air.  This  is  also  formed 
by  treating  a  heated  mixture  of  sodium  chloride  and  the 
alloy  of  osmium  and  iridium  with  chlorine  and  water 
vapor.  It  is  commonly  called  osmic  acid,  though  its  acid 
properties  are  very  weak.  Like  ruthenium  peroxide  it 
is  volatile.  It  sublimes  in  colorless,  lustrous  needles, 
and  boils  without  decomposition  at  a  temperature  a  little 
above  100°.  It  has  an  intense  odor  similar  to  that  of 
chlorine,  and  its  vapor  attacks  the  eyes  and  respiratory 
organs  somewhat  in  the  same  way  that  chlorine  does. 
It  dissolves  slowly  in  water,  and  reducing  agents  pre- 
cipitate the  metal  from  the  solution.  A  solution  of  osmic 
acid  is  used  in  microscopic  work.  When  injected  into 
the  tissues,  the  parts  are  hardened  and  colored.— 
Potassium  osmite,  K2OsO4,  is  formed  when  a  solution  of 
the  peroxide  in  potassium  hydroxide  is  treated  with  a  re- 
ducing agent.  It  is  easily  decomposed  in  water  solution. 
— The  oxides  OsO,  OsO2,  and  Os2O3  have  neither  acid 
nor  basic  properties. 

BHODIUM,  Eh  (At.  Wt.  102.23). 

Ehodium  has  no  acid  properties,  and  does  not  form  a 
peroxide  corresponding  to  those  of  ruthenium  and  os- 
mium. On  the  other  hand,  its  oxide,  Eh2O3,  is  basic. 
The  chloride  EhCl3  is  readily  formed,  and  it  is  doubt- 
ful whether  the  di-  and  tetrachlorides  have  been  made. 

IRIDIUM,  Ir  (At.  Wt.  191.66). 

Preparation. — The  extraction  of  iridium  with  plati- 
num and  with  osmium  from  platinum -ore  was  referred 
to  above.  In  order  to  separate  it  from  platinum,  advan- 
tage is  taken  of  the  fact  that  it  forms  a  trichloride,  IrCl3, 


IRIDIUM.  725 

which  with  ammonium  chloride  gives  an  easily  soluble 
double  chloride.  The  reduction  is  accomplished  either 
by  heating  the  tetrachloride  for  some  time  at  150°,  or  by 
treating  the  insoluble  double  chloride  in  water  with 
hydrogen  sulphide  or  with  sulphur  dioxide.  From  os- 
mium it  is  separated  by  treating  with  moist  chlorine, 
when,  as  stated  above,  the  osmium  is  converted  into  the 
peroxide,  which  being  volatile  passes  over.  The  residue 
contains  the  iridium  in  the  form  of  the  tetrachloride,  and 
this,  when  treated  with  potassium  chloride,  forms  the 
difficulty  soluble  chloriridate,  K2IrCl6. 

Properties. — Iridium  has  a  grayish- white  color,  and 
resembles  polished  steel.  Its  specific  gravity  is  nearly 
the  same  as  that  of  osmium,  being  22.42.  It  is  harder 
and  more  brittle  than  platinum  ;  melts  at  a  higher  tem- 
perature ;  and  unless  it  is  finely  divided  it  is  not  dis- 
solved by  nitro-hydrochloric  acid.  When  heated  with 
potassium  hydroxide  and  saltpeter  it  is  converted  into 
the  oxide. 

Chlorides. — When  finely  divided  iridium  is  treated 
with  nitro-hydrochloric  acid  it  is  converted  into  the 
tetrachloride,  IrCl4.  When  the  solution  of  the  tetra- 
chloride is  heated  it  gives  off  chlorine,  and  the  dicHoride, 
IrCl2,  is  formed.  The  trichloride,  IrCl3,  is  formed  when 
the  metal  is  heated  in  chlorine  gas.  Both  the  tetrachlo- 
ride and  the  trichloride  form  double  salts  with  the  chlo- 
rides of  the  alkali  metals.  Those  with  the  tetrachloride 
have  the  general  formula  M2IrClB,  or  IrCl4.2MCl ;  while 
those  with  the  trichloride  have  the  general  formula 
M3IrCl3,  or  IrCl,.3MCl.  The  latter  are  all  soluble  in 
water  ;  of  the  former,  the  potassium  salt,  K2IrCl,,  and 
the  ammonium  salt,  (NH4)2IrCl6,  are  almost  insoluble  in 
water. 

Oxides. — The  oxides  have  neither  acid  nor  basic  prop- 
erties. The  one  most  easily  obtained  is  the  dioxide 
IrO2.  The  hydroxides,  Ir(OH)3  and  Ir(OH)4,  are  ob- 
tained, the  former  as  a  black  and  the  latter  as  a  blue 
precipitate,  by  treating  the  chlorides  with  potassium 
hydroxide. 


726  INORGANIC  CHEMISTRY. 

PALLADIUM,  Pd  (At.  Wt.  105.56). 

Preparation. — The  chief  source  of  palladium  is  a 
Brazilian  gold-ore.  From  this  ore  the  metal  can  be  ob- 
tained by  various  methods,  one  of  which  consists  in 
melting  it  together  with  silver,  and  then  treating  it  with 
nitric  acid,  when  the  silver  and  palladium  dissolve,  and 
the  gold  remains  undissolved.  The  silver  is  precipitated 
as  chloride  and  the  palladium  as  the  cyanide,  and  when 
the  latter  is  ignited  it  is  decomposed,  leaving  palladium. 

Properties. — Palladium  resembles  iridium  and  plati- 
num in  appearance.  Its  specific  gravity  is  only  about 
half  as  great  as  that  of  platinum,  being  11.5 ;  it  is  more 
easily  fusible  than  platinum,  and  dissolves  in  nitric  acid 
and  in  hot  concentrated  sulphuric  acid.  The  property 
of  palladium  which  has  perhaps  attracted  most  atten- 
tion is  its  power  to  absorb  hydrogen,  and  form — 

Palladium-Hydrogen. — The  formation  of  this  com- 
pound was  referred  to  under  Hydrogen  (which  see).  The 
combination  takes  place  even  at  the  ordinary  tempera- 
ture, but  best  at  100°.  If  the  me-tal  is  brought  into 
hydrogen  at  this  temperature,  it  absorbs  more  than  900 
times  its  volume,  forming  an  alloy  of  the  composition 
Pd2H.  This  alloy  has  a  greater  volume  and  lower  spe- 
cific gravity  than  the  palladium  from  which  it  is  formed. 
At  130°  it  begins  to  decompose  under  the  atmospheric 
pressure,  but  continued  heating  at  a  red  heat  is  neces- 
sary to  decompose  it  completely.  If  allowed  to  lie  in 
contact  with  the  air  the  hydrogen  is  oxidized  to  water. 
Palladium-hydrogen  acts  as  a  strong  reducing  agent, 
the  hydrogen  which  it  gives  up  being  apparently  in  the 
nascent  or  atomic  condition. 

Chlorides. — When  palladium  is  dissolved  in  concen- 
trated nitro-hydrochloric  acid  it  is  converted  into  palladic 
chloride,  PdCl4,  which  with  the  chlorides  of  the  alkali 
metals  forms  double  salts  similar  to  those  formed  by 
iridium  tetrachloride,  and,  as  we  shall  see,  by  platinic 
chloride.  The  tetrachloride  is  decomposed  by  evapo- 
ration of  its  solution,  giving  up  chlorine  and  leaving  pal- 
ladious  chloride,  PdCla,  which  crystallizes,  and  forms 


PLATINUM.  727 

with  the  chlorides  of  the  alkali  metals  double  salts  of  the 
general  formula  M2PdCl4,  or  PdCl2.2MCl. 

Oxides. — The  point  of  chief  interest  presented  by  the 
oxides  is  that  in  one  of  them,  the  suboxide,  Pd2O,  the 
metal  appears  as  a  univalent  element.  The  dioxide  or 
palladia  oxide,  PdO2,  has  neither  acid  nor  basic  proper- 
ties. The  monoxide,  or  palladious  oxide,  PdO,  forms 
unstable  salts  with  acids,  an  example  being  the  sulphate, 
PdS04  +  2HaO. 

PLATINUM,  Pt  (At.  Wt.  193.41). 

Preparation. — A  general  idea  of  the  method  of  pro- 
cedure in  extracting  platinum  from  its  ores  was  given 
on  p.  722.  Thus  prepared,  however,  it  always  contains 
iridium,  and  for  some  purposes  for  which  platinum  is 
used  this  is  objectionable.  In  order  to  purify  the  metal 
advantage  is  taken  of  the  fact  that  iridium  chloride  can 
be  converted  into  a  trichloride,  which  with  ammonium 
chloride  forms  an  easily  soluble  double  salt  (see  p. 
725).  The  metal  as  obtained  by  igniting  ammonium 
platinic  chloride  forms  a  gray  spongy  mass  known  as 
spongy  platinum.  When  a  solution  of  platinous  chloride 
is  boiled  with  potassium  hydroxide,  and  alcohol  gradually 
added,  the  salt  is  reduced,  and  the  platinum  is  precipitated 
as  an  extremely  fine  powder,  known  as  platinum  black. 
When  spongy  platinum  and  platinum  black  are  heated 
to  fusion  by  the  oxyhydrogen  flame  they  are  converted 
into  the  compact  variety. 

Properties. — Platinum  is  a  grayish-white  metal  re- 
sembling polished  steel ;  it  can  be  drawn  out  into  very 
fine  wire  ;  it  melts  in  the  flame  of  the  oxyhydrogen  blow- 
pipe, and  when  heated  above  its  melting-point  it  is 
volatile ;  its  specific  gravity  is  21.5.  At  white  heat  it 
can  be  welded.  It  is  not  dissolved  by  nitric  acid,  hydro- 
" chloric  acid,  nor  sulphuric  acid,  but  it  dissolves  in  nitro- 
hydrochloric  acid,  forming  the  acid,  H2PtCl6.  Fusing 
alkalies,  and  particularly  a  mixture  of  caustic  potash 
'and  saltpeter,  act  upon  it;  but  the  alkaline  carbonates 
do  not.  In  contact  with  red-hot  charcoal  and  silicon 
dioxide  a  compound  of  silicon  and  platinum  is  formed. 


728  INORGANIC  CHEMISTRY. 

Finely  divided  platinum  has  to  a  remarkable  extent  the 
power  of  condensing  gases  upon  its  surface.  It  ab- 
sorbs, for  example,  200  times  its  own  volume  of  oxy- 
gen, and  other  gases  in  a  similar  way.  The  oxygen 
thus  absorbed  is  in  active  condition,  and  if  oxidizable 
substances  are  brought  in  contact  with  it  they  are  easily 
oxidized.  Thus  when  a  current  of  hydrogen  is  allowed 
to  flow  against  a  piece  of  spongy  platinum  it  takes  fire, 
owing  to  the  presence  of  the  condensed  oxygen  in  the 
pores  of  the  platinum.  Similarly,  when  sulphur  dioxide 
and  oxygen  are  allowed  to  flow  together  over  spongy 
platinum,  or  even  the  compact  metal,  the  two  gases  unite 
to  form  sulphur  trioxide. 

Applications  of  Platinum. — The  metal  is  of  the  great- 
est value  to  the  chemist  on  account  of  its  power  to  resist 
the  action  of  high  temperatures  and  of  most  chemical 
substances.  It  is  used  in  the  laboratory  in  the  form  of 
wire,  foil,  crucibles,  evaporating-dishes,  tubes,  etc.,  etc. 
From  what  was  said  above  it  cannot  be  used  with 
alkalies  and  saltpeter,  nor  with  nitro-hydrochloric  acid. 
Platinum  vessels,  further,  should  not  be  placed  upon 
red-hot  charcoal.  Metallic  salts  which  are  easily  re- 
duced, such  as  those  of  antimony  and  bismuth,  should 
not  be  heated  in  platinum  vessels,  as  the  reduced  ele- 
ments, like  silicon,  form  alloys  with  the  platinum,  and 
these,  as  a  rule,  are  easily  fusible.  In  the  concentration 
of  sulphuric  acid  on  the  large  scale  platinum  stills  are 
used.  The  price  of  platinum  is  not  as  high  as  that  of 
gold,  but  much  higher  than  that  of  silver. 

Alloys  of  Platinum. — The  only  alloy  of  platinum  which 
is  of  any  special  importance  is  that  which  it  forms  with 
iridium.  A  small  percentage  of  iridium  diminishes  the 
malleability  of  platinum  very  markedly,  and  makes  it 
brittle  ;  it,  however,  increases  its  resistance  to  the  action 
of  reagents.  An  alloy  of  90  per  cent  platinum  and  10 
per  cent  iridium  has  been  adopted  by  the  French  Gov- 
ernment as  the  best  material  from  which  to  make  normal 
meters.  This  alloy  is  very  hard,  as  elastic  as  steel,  more 
difficultly  fusible  than  platinum,  entirely  unchangeable 
in  the  air,  and  is  capable  of  a  high  polish. 


COMPOUNDS  OF  PLATINUM.  729 

Chlorides.  —  Like  palladium,  platinum  forms  two  chlo- 
rides, platinom  chloride,  PtCl2,  and  platinic  chloride,  PtCl4. 
The  latter  is  formed  when  platinum  is  dissolved  in  aqua 
regia,  and  the  solution  evaporated  to  dryness.  From  its 
solution  in  water  it  crystallizes  with  ten  or  five  molecules 
of  water.  It  is  soluble  in  alcohol  as  well  as  in  water. 
When  the  dry  substance  is  heated  for  some  time  to 
225°-230°,  it  is  decomposed,  yielding  platinous  chloride, 
which  is  a  grayish-green  powder  insoluble  in  water. 

Chlorplatinic  Acid,  H2PtCl6,  is  formed  by  direct  union 
of  platinic  chloride  with  hydrochloric  acid.  It  crystal- 
lizes with  six  molecules  of  water,  and  forms  a  series  of 
salts  called  the  chlorplatinates,  to  which  reference  has 
already  been  made.  Those  most  commonly  met  with 
in  the  laboratory  are  the  potassium  salt,  K2PtCl6,  or 
PtCl4.2KCl,  and  the  ammonium  salt,  (NH4)2PtCl6,  or 
PtCl4.2NH4Cl,  both  of  which  are  difficultly  soluble  in 
water,  and  are  therefore  precipitated  when  platinic  chlo- 
ride is  added  to  solutions  containing  potassium  or  ammo- 
nium chloride.  The  sodium  salt  is  easily  soluble  in  water. 
Many  other  chlorplatinates  are  known,  and  many  crys- 
tallize well.  Considering  the  similarity  in  composition  be- 
tween chlorplatinic  acid  and  fluosilicic  acid,  the  conclu- 
sion seems  justified  that  they  have  the  same  constitution. 
The  reasons  which  lead  to  the  belief  that  the  constitution 
of  fluosilicic  acid  is  properly  represented  by  the  formula 
fFl 

I     T^l 

Si  4  ™  v  jj  make  it  probable  that  the  constitution  of 
chlorplatinic  acid  should  be  represented  by  a  similar 
formula,  Pt 


Platinous  chloride  like  platinic  chloride  combines  with 
other  chlorides  to  form  double  salts,  the  general  formula 
of  which  is  M2PtCl4,  or  PtCl2.2MCl. 

Cyanides.  —  Platinum  forms  a  number  of  beautiful 
double  cyanides  derived  from  an  acid  of  the  formula 


730  INORGANIC  CHEMISTRY. 

H2Pt(CN)4  or  Pt(CN)2.2HCN,  which  should  be  called 
cyanplatinous  acid.  It  is  analogous  to  the  acid  from 
which  the  double  chlorides  of  platinous  chloride  are 
derived,  H2PtCl4.  These  cyanplatinites  are  easily  ob- 
tained, and,  as  a  rule,  crystallize  well  and  are  beauti- 
fully colored.  The  magnesium  salt,  MgPt(CN)4  +  7H2O, 
forms  quadratic  prisms,  the  side  faces  of  which  have  a 
green  metallic  lustre,  while  the  end  faces  are  deep  blue. 
Hydroxides  and  Oxides. — When  a  solution  of  platinic 
chloride  is  treated  with  sodium  hydroxide,  and  afterward 
with  acetic  acid,  a  white  precipitate  of  platinic  hydroxide, 
Pt(OH)4  +  2H2O,  is  formed,  which  when  dried  at  100° 
loses  water  and  is  converted  into  the  brown  hydroxide, 
Pt(OH)4.  This  loses  water  when  heated  higher  and  is 
converted  into  the  oxide,  PtO2.  In  a  similar  way  plati- 
nous hydroxide,  Pt(OH)2,  and  platinous  oxide,  PtO,  are 
obtained  from  platinous  chloride.  Platinic  hydroxide,, 
Pt(OH)4,  has  acid  properties,  and  forms  a  few  salts  of 
the  general  formula  M2PtO3,  of  which  barium  platinate, 
BaPtO3,  is  the  best  known.  Platinic  acid,  from  which 
these  salts  are  derived,  is  plainly  formed  from  platinic  hy- 
droxide by  loss  of  one  molecule  of  water,  and  bears  to  it 
the  same  relation  that  ordinary  silicic  acid,  H2SiO3,  bears 
to  normal  silicic  acid,  Si(OH)4.  Further,  platinic  acid 
and  chlorplatinic  acid  appear  to  be  analogous  compounds; 
and  the  latter  may  be  regarded  as  derived  from  the 
former  by  replacement  of  the  three  atoms  of  oxygen 
by  six  atoms  of  chlorine,  as  shown  in  the  formulas  ] 

H2Pt03; 


Sulphides. — There  are  two  sulphides  of  platinum  which 
are  analogous  to  the  two  oxides,  PtO  and  PtO2.  These 
are  platinous  sulphide,  PtS,  and  platinic  sulphide,  PtS2> 
They  are  black  insoluble  compounds,  which  are  precip- 
itated when  hydrogen  sulphide  or  soluble  sulphides  are 
added  to  solutions  of  platinous  and  platinic  chlorides. 


THE  PLATINUM  BASES.  731 

Compounds  with  Ammonia — The  Platinum  Bases. — Like 
cobalt  salts,  the  salts  of  platinum  form  a  large  number 
of  compounds  with  ammonia.  When  ammonia  acts  upon 
a  solution  of  platinous  chloride  a  compound  of  the  for- 
mula PtCl2(NH3)2  is  formed.  This  is  the  starting-point 
for  a  series  of  compounds,  as  the  bromide,  PtBr2(NH3)2 ; 
the  nitrate,  Pt(NO3)2(NH3)2 ;  the  sulphate,  PtSO4(NH3)2 ; 
etc.  There  is  another  series  beginning  with  the  chlo- 
ride, PtCl2(NH3)3 ;  another  beginning  with  the  chloride, 
PtCl2(NH3)4.  All  the  above  are  to  be  regarded  as  derived 
from  platinous  chloride.  Similarly  there  are  other  series 
obtained  from  platinic  chloride.  The  chlorides  have 
the  formulas  PtCl4(NH3)2,  PtCl4(NH3)3,  PtCl4(NH3)4.  It 
seems  probable  that  these  salts  are  ammonium  salts 
in  which  a  part  of  the  hydrogen  of  the  ammonium  is 
replaced  by  platinum.  Thus  the  chloride  PtCla(NH8), 

probably  has  the  constitution  P^<NH8fT     Although  a 

great  deal  of  work  has  been  done  on  these  platino- 
ammonium  compounds,  and  much  interesting  informa- 
tion in  regard  to  them  has  been  gained,  the  subject  of 
their  constitution  is  still  in  an  unsatisfactory  state. 


APPENDIX  T 

CONTAINING  SPECIAL  DIRECTIONS  FOR 
LABORATORY   WORK. 

Introduction. — In  order  to  become  familiar  with  the  prin- 
ciples of  Chemistry  it  is  absolutely  necessary  that  the  student 
should  devote  a  part  of  his  time  to  work  in  the  laboratory — > 
the  more  the  better.  It  is,  further,  necessary  that  the  labora- 
tory work  should  be  done  with  the  greatest  care.  Every  piece 
of  apparatus  should  be  carefully  constructed,  the  desk  should 
be  kept  clean  and  in  good  order;  and  no  work  should  be 
abandoned  until  the  student  is  satisfied  that  he  has  seen  all 
there  is  to  be  seen,  and  that  he  has  learned  all  that  the  work 
can  teach  him.  He  must  learn  to  use  his  own  senses,  and  to 
believe  what  he  sees,  and  not  simply  "what  the  book  says." 
It  sometimes  happens  that  owing  to  the  peculiar  way  in  which 
an  experiment  is  performed  results  quite  different  from  those 
anticipated  are  obtained.  Under  these  circumstances  it  is  not 
advisable  to  conclude  at  once  that  "  the  book  must  be  wrong." 
It  may  be;  but  the  probabilities  are  against  this  explanation 
of  the  discrepancy.  Nothing  is  more  instructive  than  well- 
directed  efforts  to  find  the  causes  of  difficulties.  Such  efforts, 
more  than  anything  else,  develop  the  spirit  of  true  scientific 
inquiry.  It  is  advisable  for  the  student  to  carry  on  the  work 
for  which  directions  are  given  below  in  connection  with  the 
study  of  the  book.  A  good  plan  to  follow  is  to  read  a  chapter 
with  care;  then  to  perform  the  experiments  which  are  intended 
to  illustrate  that  chapter,  and,  while  doing  the  work,  again  to 
read;  and  afterwards  to  write  out  an  account  of  what  has  been 
done,  noting  everything  of  importance  exactly  as  it  was  ob- 
served. If  experiments  are  necessary  to  account  for  phe- 
nomena not  described  in  the  book,  these  should  be  described; 
and  if  a  conclusion  is  reached  in  regard  to  these  phenomena, 
the  evidence  upon  which  the  conclusion  is  based  should  be 
clearly  stated.  It  is  only  by  patient  work  carried  on  in  this 
way  that  one  can  hope  to  reach  a  clear  conception  of  the 
science.  But  by  such  work  the  desired  result  will  be  reached. 

(733) 


734       EXPERIMENTS  TO  ACCOMPANY  CHAPTER  I. 

Progress  will  seem  slow  at  first,  as  it  always  does  in  a  new  sub- 
ject; but  in  time  the  ideas  will  begin  to  arrange  themselves 
systematically,  order  will  come  out  of  confusion,  and  in  this 
result  the  conscientious  student  will  find  a  delightful  reward 
for  his  labor.  Every  great  branch  of  knowledge  is  made  up 
of  details  which  are  bound  together  by  certain  broad  govern- 
ing principles.  It  is  impossible  to  avoid  these  details.  In 
order  to  understand  the  governing  principles  the  details 
must  be  studied  to  some  extent.  They  form  the  raw  material 
from  which  the  science  is  constructed;  without  them  the 
science  would  be  impossible.  As  well  might  one  hope  to  learn 
a  language  by  studying  its  grammatical  rules  and  avoiding  the 
details  of  the  mere  words,  as  to  learn  chemistry  by  studying 
the  laws  and  avoiding  contact  with  the  things  to  which  these 
laws  have  reference. 


EXPERIMENTS  TO  ACCOMPANY   CHAPTER  I. 
CHEMICAL  CHANGE  CAUSED  BY  HEAT. 

Experiment  1 — In  a  clean  dry  test-tube  put  enough  white 
sugar  to  make  a  layer  £  to  £  an  inch  thick.  Hold  the  tube 
in  the  flame  of  a  spirit-lamp  or  a  laboratory  burner.  What 
evidence  is  furnished  by  this  experiment  that  chemical  change 
may  be  caused  by  heat?  What  is  left  in  the  tube?  Is  it 
soluble?  Is  it  sweet  ?  Is  it  sugar? 

Experiment  2. — Half-fill  the  bulb  of  an  arsenic-tube  with 
red  oxide  of  mercury,  or,  if  such  a  tube  is  not  available,  pro- 
ceed as  follows:  From  a  piece  of  hard-glass  tubing  of  about 
6  to  7  millimeters  (^  inch)  internal  diameter  cut  off  a  piece 
about  10  centimeters  (4  inches)  long  by  making  a  mark 
across  it  with  a  triangular  file,  and  then  seizing  it  with  both 
hands,  one  on  each  side  of  the  mark,  pulling  and  at  the 
same  time  pressing  slightly  as  if  to  break  it.  Clean  and  dry 
it,  and  hold  one  end  in  the  flame  of  a  blast-lamp  until  it 
melts  together.  During  the  melting  turn  the  tube  constantly 
around  its  long  axis  so  that  the  heat  may  act  uniformly  upon 
it.  Put  into  the  tube  thus  made  enough  red  oxide  of  mer- 
cury (mercuric  oxide)  to  form  a  layer  about  12  millimeters 
(i  inch)  thick.  Heat  the  tube  as  in  the  last  experiment. 
What  change  in  color  is  noticed  ?  What  is  deposited  upon 
the  glass  in  the  upper  part  of  the  tube  ?  What  evidence  is 


CHANGES  EFFECTED  BY  AN  ELECTRIC  CURRENT.    735 


furnished  by  this  experiment  that  chemical  change  can  be 
effected  by  heat  ? 

CHEMICAL  CHANGES  CAN  BE   EFFECTED  BY  AN  ELECTRIC 

CURRENT. 

Experiment  3. — To  the  ends  of  insulated  copper  wires  con- 
nected with  two  cells  of  a  Bunsen's  or  Grove's  battery  fasten 
platinum  plates,  say  25  mm.  (1  inch)  long  by  12  mm.  (£  inch) 
wide.  Insert  these  platinum  electrodes  into  water  contained 
in  a  shallow  glass  vessel  about  15  cm.  (6  inches)  wide  and  7 
to  8  cm.  (3  inches)  deep,  taking  care  to  keep  them  separated 
from  each  other.  No  action  will  take  place,  for  the  reason, 
as  has  been  shown,  that  water  will  not  conduct  the  current, 
and  hence  when  the  platinum  electrodes  are  kept  apart  there 
is  no  cucrent.  By  adding  to  the  water  about  one  tenth  its 
own  volume  of  strong  sulphuric  acid  it  acquires  the  power  to 
convey  the  current.  It  will  then  be  observed  that  bubbles 
rise  from  each  of  the  platinum  plates.  In  order  to  collect 
them  an  apparatus  like  that  shown  in  Fig.  15  may  be  used. 

h  and  o  represent  glass  tubes  which  may  conveniently  be  about 
30  cm.  (1  foot)  long  and  25  mm.  (1  inch)  internal  diameter. 
They  are  first  filled  with  the  water  containing  one  tenth  its  vol- 
ume of  sulphuric  acid,  and  then  placed  with  the  mouth  under 
water  in  the  vessel  J.  The  platinum  elec- 
trodes are  now  brought  beneath  the  invert- 
ed tubes.  The  b  ubbles  which  rise  from  them 
will  pass  upward  in  the  tubes  and  the  water 
will  be  pressed  down.  Gradually  the  water 
will  be  completely  forced  out  of  one  of  the 
tubes,  while  the  other  is  still  half  full  of  wa- 
ter. The  substances  thus  collected  in  the 
tubes  are  invisible  gases.  After  the  first  tube 
is  full  of  gas,  place  the  thumb  over  its  mouth 
and  remove  the  tube.  Turn  it  mouth  up- 
ward, and  at  once  apply  a  lighted  match  to  it. 
A  flame  will  be  noticed.  The  gas  which  was 
contained  in  the  tube  is  therefore  capable  of 
burning.  It  cannot,  therefore,  have  been 
air.  In  the  mean  time  the  second  tube  will 
have  become  filled  with  gas.  Kemove  this  tube  in  the  same 
way  and  insert  a  thin  piece  of  wood  with  a  spark  on  it.  The 
spark  will  at  once  burst  into  flame,  and  the  burning  of  the 


FIG.  15. 


736       EXPERIMENTS  TO  ACCOMPANY  CHAPTER  I. 

wood  will  take  place  more  actively  than  it  does  in  ordinary  airy 
as  may  be  shown  by  withdrawing  it  and  again  inserting  it  inta 
the  tube.  The  gas  in  this  tube,  it  will  be  noticed,  does  not 
take  fire.  Without  going  into  further  details,  it  is  clear  from 
the  above  experiment  that  when  an  electric  current  acts  on 
water  two  invisible  gases  are  produced.  A  chemical  change  is 
caused  by  an  electric  current. 

MECHANICAL  MIXTURES  AND  CHEMICAL  COMPOUNDS. 

Experiment  4. — Examine  carefully  a  piece  of  coarse-grained 
granite;  break  off  some  of  it,  and  separate  the  constituents. 
How  many  are  there?  By  what  properties  do  you  recognize 
them  ?  Powder  a  small  bit  of  one  of  the  constituents,  and 
examine  the  powder  with  the  microscope.  Do  you  recognize 
more  than  one  kind  of  matter  ?  Mix  the  powder  of  the  three 
constituents,  and  see  whether  in  the  mixed  powder  there  is 
any  difficulty  in  detecting  the  three  kinds  of  matter  with  the 
aid  of  the  microscope. 

Experiment  5. — Mix  a  gram  or  two  of  powdered  roll- 
sulphur  and  an  equal  weight  of  very  fine  iron  filings  in  a 
small  mortar.  Examine  a  little  of  the  mixture  with  a  micro- 
scope. Not  only  can  we  recognize  the  particles  of  iron  and  of 
sulphur  by  means  of  the  microscope,  but  we  can  also  pick  out 
the  pieces  of  iron  by  means  of  a  magnet.  The  magnet  attracts 
the  iron  but  not  the  sulphur,  so  that  by  passing  the  magnet 
often  enough  through  the  mixture  we  can  pick  out  all  the 
iron  and  leave  all  the  sulphur.  This  separation  is  really  a 
mechanical  separation.  It  is  only  a  somewhat  more  refined 
method  of  picking  out  than  that  used  in  the  case  of  granite. 

Experiment  6. — Pass  a  small  magnet  through  the  mixture 
above  prepared.  Unless  the  substances  used  are  thoroughly 
dry,  particles  of  sulphur  will  adhere  to  the  magnet,  but  even 
then  it  will  be  seen  that  most  of  that  which  is  taken  out  of 
the  mixture  is  iron. 

The  iron  and  sulphur  can  also  be  separated  by  treating  the 
mixture  with  a  liquid  known  as  carbon  disulphide.  Sulphur 
dissolves  in  this  liquid,  but  iron  does  not.  So  that  when  the 
mixture  is  treated  with  it  the-  iron  is  left  behind,  and  can 
easily  be  recognized  as  such. 

Experiment  7. — Pour  two  or  three  cubic  centimeters  of 
carbon  disulphide  on  a  little  powdered  roll-sulphur  in  a  dry 
test-tube.  The  sulphur  dissolves.  Treat  iron  filings  in  the 


VARIOUS  EXAMPLES  OF  CHEMICAL  ACTION.       737 

same  way.  The  iron  does  not  dissolve.  Now  treat  a  small 
quantity  of  the  mixture  with  carbon  disulphide.  After  the 
sulphur  is  dissolved  pour  off  the  solution  on  a  good-sized 
watch-glass  and  let  it  stand.  Examine  what  remains  undis- 
solved  in  the  test-tube,  and  satisfy  yourself  that  it  is  iron. 
After  the  liquid  has  evaporated  examine  what  is  left  in  the 
watch-glass  and  satisfy  yourself  that  it  is  sulphur.  Why  are 
you  justified  in  concluding  that  the  substance  left  in  the  test- 
tube  is  iron  and  that  left  on  the  watch-glass  is  sulphur  ? 

Experiment  8. — Make  a  fresh  mixture  of  three  grams  each 
of  powdered  roll-sulphur  and  fine  iron  filings.  Grind  them 
together  intimately  in  a  dry  mortar  and  put  them  in  a  dry 
test-tube.  Heat  gradually  until  the  mass  begins  to  glow.  At 
first  the  sulphur  melts  and  becomes  dark-colored.  It  may 
even  take  fire.  But  soon  something  else  evidently  takes  place. 
The  whole  mass  begins  to  glow,  and  if  you  at  once  take  the 
tube  out  of  the  flame,  the  mass  continues  to  glow,  becoming 
brighter.  This  soon  stops;  the  mass  grows  dark  and  gradually 
cools  down.  As  soon  as  it  reaches  the  ordinary  temperature, 
the  tube  should  be  broken  and  the  contents  put  in  a  mortar. 
A  close  examination  will  show  that  the  mass  does  not  look 
like  the  mixture  of  sulphur  and  iron  with  which  we  started. 
It  has  a  bluish-black  color,  and  is  apparently  homogeneous. 
An  examination  with  the  microscope,  the  magnet,  and  car- 
bon disulphide  will  prove  that,  while  there  may  be  a  little 
iron  left,  and  possibly  a  little  sulphur,  most  of  the  bluish- 
black  mass  is  neither  iron  nor  sulphur,  but  a  new  substance 
with  properties  quite  different  from  those  of  iron  and  from 
those  of  sulphur. 

OTHER  EXAMPLES  OF  CHEMICAL  ACTION. 

Experiment  9. — Examine  a  piece  of  calc-spar  or  marble. 
You  see  that  it  is  made  up  of  pieces  of  definite  shape.  It  is, 
as  we  say,  crystallized.  It  is  quite  hard,  though  a  knife  will 
cut  it.  Heated  in  a  hard-glass  tube,  as  in  Experiment  2,  it 
does  not  melt,  but  remains  essentially  unchanged.  It  does 
not  dissolve  in  water.  To  prove  this,  put  a  piece  the  size  of  a 
pea  in  a  test-tube  with  pure  water.  Thoroughly  shake,  and 
then,  as  heating  usually  aids  solution,  boil.  Now  pour  off  a 
few  drops  of  the  liquid  on  a  piece  of  platinum-foil  or  a  watch- 
glass,  and  by  gently  heating  cause  the  water  to  evaporate.  If 
there  is  anything  in  solution,  there  will  be  a  solid  residue  on 


738       EXPERIMENTS  TO  ACCOMPANY  CHAPTER  1. 

the  platinum-foil  or  watch-glass.  If  not,  there  will  be  no 
residue.  Now  treat  a  small  piece  of  the  substance  with  dilute 
hydrochloric  acid  and  notice  what  takes  place.  Bubbles  of  gas 
are  given  off.  After  the  action  has  continued  for  about  a 
minute,  insert  a  lighted  match  in  the  upper  part  of  the  tube. 
It  is  extinguished,  and  the  gas  does  not  burn.  The  gas  formed 
in  this  case  is  therefore  plainly  not  identical  with  either  one 
of  those  obtained  from  water  b}^  the  action  of  the  electric  cur- 
rent (see  Experiment  3).  It  is  what  is  commonly  called  car- 
bonic-acid gas.  As  the  action  continues,  the  piece  of  calc-spar 
or  marble  grows  smaller  and  smaller,  and  finally  disappears, 
when  there  is  a  clear  solution.  The  substance  has  dissolved 
in  the  hydrochloric  acid.  In  order  to  determine  whether  any- 
thing else  has  taken  place  besides  the  dissolving,  we  shall  have 
to  get  rid  01  the  excess  of  hydrochloric  acid.  This  we  can 
easily  do  by  boiling  it,  when  it  passes  off  in  the  form  of  vapor, 
and  then  whatever  is  in  solution  will  remain  behind.  For  this 
purpose  put  the  solution  in  a  small,  clean  porcelain  evaporat- 
ing-dish,  and  put  this  on  a  vessel  containing  boiling  water,  or 
a  water-bath.  The  operation  should  be  carried  on  in  a  place 
in  which  the  draught  is  good,  so  that  the  vapors  will  not  collect 
in  the  working-room.  They  are  not  poisonous,  but  they  are 
annoying.  The  arrangement  for  evaporating  is  represented  in 
Fig.  16. 

After  the  liquid  has  evaporated  and  the  suoatance  in  the 

evaporating-dish  is  dry,  examine 
it,  and  carefully  compare  its  prop- 
erties with  those  of  the  substance 
which  was  put  into  the  test-tube. 
Its  structure  wilj  be  found  not  to 
present  the  regularities  noticed  in 
the  original  substance.  It  is  much 
softer,  dissolv(nS  in  water,  melts 
when  heated  in  a  hard-glass  tube. 
It  does  not  give  off  a  gas  when 
treated  with  hydrochloric  acid. 
When  exposed  to  the  air  it  soon 

FIG.  16.  ,  . 

becomes  moist,  and  after  a  time 

liquid.  The  experiment  shows  that  when  hydrochloric  acid 
acts  upon  calc-spar  or  marble  the  latter  at  least  loses  its  own 
properties.  It  might  be  shown  that  some  of  the  hydrochloric 
acid  also  loses  its  properties.  In  place  of  tha  two  wo  get  a  new 


VARIOUS  EXAMPLES  OF  CHEMICAL  ACTION.        739 

substance  with  entirely  different  properties.  The  two  sub- 
stances have  acted  chemically  upon  each  other  arid  produced, 
a  chemical  compound.  In  this  case  it  was  only  necessary  to 
bring  the  substances  in  contact  in  order  to  cause  them  to  act 
chemically  upon  each  other.  It  was  not  necessary  to  heat 
them,  as  it  was  in  the  case  of  the  iron  and  sulphur. 

Experiment  1O. — Bring  together  in  a  test-tube  a  small 
piece  of  copper  and  some  moderately  dilute  nitric  acid.  In  a 
short  time  action  begins.  The  upper  part  of  the  tube  becomes 
filled  with  a  dark,  reddish-brown  gas  which  has  a  disagreeable 
smell.  Do  not  inhale  it,  as  when  taken  into  the  lungs  it 
produces  bad  effects.  The  solution  becomes  colored  dark 
blue,  and  the  copper  disappears.  Examine  this  solution, 
as  in  Experiment  9,  and  see  what  has  been  formed.  What  are 
the  properties  of  the  substance  found  after  evaporation  of  the 
liquid?  Is  it  colored?  Is  it  soluble  in  water?  Does  it  change 
when  heated  in  a  tube  ?  Is  it  hard  or  soft  ?  Does  it  in  any 
way  suggest  the  copper  with  which  you  started  ? 

Experiment  11. — Try  the  action  of  dilute  sulphuric  acid 
on  a  little  zinc  in  a  test-tube.  A  gas  will  be  given  off.  Apply 
a  lighted  match  to  it.  Does  the  result  suggest  anything  no- 
ticed in  an  experiment  already  performed?  After  the  zinc 
has  disappeared,  evaporate  the  solution  as  in  Experiment  9. 
Carefully  compare  the  properties  of  the  substance  left  behind 
with  those  of  zinc. 

Experiment  12. — Hold  the  end  of  a  piece  of  magnesium 
ribbon  about  20  centimeters  (8  inches)  long  in  a  flame  until  it 
takes  fire;  then  hold  the  burning  substance  quietly  over  a 
piece  of  dark  paper,  so  that  the  light,  white  product  may 
be  collected.  Compare  the  properties  of  this  white  product 
with  those  of  the  magnesium.  Here  again  a  chemical  act  has 
taken  place.  The  magnesium  has  combined  with  something 
which  it  found  in  the  air,  and  heat  was  produced  by  the  com- 
bination. The  product  is  the  white  substance. 

Experiment  13. — In  a  small,  dry  flask  (400  to  500  ccm.) 
put  a  bit  of  granulated  tin.  Pour  upon  it  2  or  3  ccm.  concen- 
trated nitric  acid.  If  no  change  takes  place,  heat  gently,  and 
presently  there  will  be  a  copious  evolution  of  a  reddish-brown 
gas  with  a  disagreeable  smetl,  (under  what  conditions  has  a 
gas  like  this  already  been  obtained  ?)  the  tin  will  disappear, 
and  in  its  place  will  appear  a  white  powder.  Compare  the 


740       EXPERIMENTS  TO  ACCOMPANY  CHAPTER  II. 

properties  of  this  white  powder  with  those  of  tin.     Why  are 
you  justified  in  concluding  that  they  are  not  the  same  thing  ? 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  II. 
PREPAKATIOX  OF  OXYGEN. 

Experiment  14.— Make  a  small  quantity  of  oxygen  by 
heating  strongly  an  arsenic-tube  half-filled  with  manganese 
dioxide  and  fitted  with  a  small  rubber  delivery-tube. 

Experiment  15. — Make  some  oxygen  by  heating  a  few 
grams  of  mercuric  oxide  in  a  hard-glass  tube  closed  at  one 
end  and  connected  at  the  other  end  by  means  of  a  cork  with 
a  bent  glass  tube. 

Experiment  16. — Arrange  an  apparatus  as  shown  in  Fig. 
17.  A  represents  a  flask  of  100  ccm.  capacity.  By  means  of 


FIG.  17. 

a  well-fitting  rubber  stopper  one  end  of  the  bent-glass  tube  B 
is  connected  with  it,  and  the  other  end,  which  should  turn 
upward  slightly,  is  placed  under  the  surface  of  the  water  in  C. 
In  A  put  4  to  5  grams  (about  an  eighth  of  an  ounce)  potas- 
sium chlorate,  and  gently  heat  by  means  of  the  lamp.  Notice 
carefully  what  takes  place.  At  first  the  potassium  chlorate 
will  melt,  forming  a  clear  liquid.  If  the  heat  is  increased, 
the  liquid  will  appear  to  boil,  and  it  will  soon  be  seen  that  a 
gas  is  given  off,  Now  bring  the  inverted  cylinder  D  filled 
with  water  over  the  end  of  the  tube,  and  let  the  bubbles 


MEASUREMENT  OF  THE  VOLUME  OF  GASES.      741 

of  gas  rise  in  the  cylinder.  After  a  considerable  quantity  of 
gas  has  been  collected  in  this  way  the  action  stops,  the  mass 
in  the  flask  becomes  solid,  and  apparently  the  end  of  the  pro- 
cess is  reached.  But  if  the  heat  is  raised  again,  gas  will  again 
begin  to  come  off,  and  in  this  second  stage  a  larger  quantity 
will  be  collected  than  in  the  first.  Finally,  however,  the  end 
is  reached,  and  the  substance  left  in  the  flask  remains  un- 
changed, no  matter  how  long  heat  may  be  applied.  An  ex- 
amination of  the  gas  collected  will  show  that  a  piece  of  wood 
will  burn  in  it  very  readily.  Explain  the  changes  which  have 
taken  place  in  this  experiment.  Calculate  how  much  oxygen 
can  be  obtained  by  heating  12  grams  of  potassium  chlorate. 

MEASUREMENT  or  THE  VOLUME  OF  GASES. 

In  studying  chemical  changes  it  often  becomes  necessary  to 
measure  the  volume  of  a  gas,  and  it  is  important  to  know  what 
precautions  must  be  taken  in  such  cases.  For  the  purpose  a 
tube  is  used  which  is  graduated  by  marks  etched  on  the  out- 
side. These  marks  may  either  indicate  the  number  of  cubic 
centimeters  of  gas  contained  in  the  tube,  or  the  length  of  the 
column  of  gas.  In  the  latter  case  it  is  of  course  necessary  to 
determine  what  volume  corresponds  to  a  given  length  of  the 
column.  The  chief  difficulty  encountered  in  measuring  gas 
volumes  is  due  to  the  fact  that  the  volume  varies  with  the 
temperature  and  pressure.  When  the  temperature  of  a  gas  is 
raised  one  degree  centigrade  its  volume  is  increased  ^-fj  part. 
If,  therefore,  the  volume  of  a  gas  at  0°  is  V,  at  t°  its  volume 
V  will  be 


This  expression  may  also  be  written  thus : 

V'  =  F-f  0.00366*  .  F,     or     V  =  F(l  +  0.003660. 
From  these  we  get  the  expressions 

97Q  V'  V 

V=^r-<    and-  F  = 


4-  t  1  +  0.00366** 

It  is  customary  to  reduce  the  observed  volume  of  a  gas  to  the 
volume  which  it  would  have  at  0°.     The  correction  is  easily 


742      EXPERIMENTS  TO  ACCOMPANY  CHAPTER  II. 

made  by  the  aid  of  the  above  formula.  Thus,  if  the  volume 
of  a  gas  is  found  to  be  250  cubic  centimeters  at  15°,  and  it  is 
required  to  know  what  the  volume  would  be  if  the  temperature 
were  reduced  to  0°,  the  calculation  is  made  thus  :  In  this 
case  the  observed  volume  V9  is  250  cc.  ;  t,  the  temperature, 
is  15°.  Substituting  these  values  in  the  equation 

273  V  V 

\/   _    _  of  I/   —    ___ 

"  273  +  t9  "  1  +  0.00366^' 

we  have 

273  X  250  250 


TF-»    or        = 


"    273  +  15  '  "  1  +  0.00366  X  15' 

from  which  we  get  236.99  as  the  value  of  F. 

But  the  volume  of  a  gas  varies  also  according  to  the  pres- 
sure. If  the  pressure  is  doubled,  the  volume  is  decreased  one 
half ;  and  if  the  pressure  is  decreased  one  half,  the  volume 
is  doubled,  and  so  on.  In  other  words,  the  volume  of  a  gas 
varies  inversely  according  to  the  pressure.  Increase  the  pres- 
sure two,  three,  or  four  times,  and  the  volume  becomes  one 
half,  one  third,  or  one  fourth,  and  vice  versa.  If  the  gas  has 
the  volume  F  at  the  pressure  P,  and  at  pressure  Pf  the  vol- 
ume V't  these  values  are  found  to  bear  to  one  another  the 
relations  expressed  in  the  equation 

VP=  V'P'. 

The  pressure  is  usually  stated  in  millimeters,  and  reference  is 
to  the  height  of  a  column  of  mercury  which  the  pressure  cor- 
responds to.  A  gas  contained  in  an  open  vessel,  or  in  a  vessel 
over  mercury  or  water,  in  which  the  level  of  the  liquid  inside 
and  outside  the  vessel  is  the  same,  is  under  the  pressure  of 
the  atmosphere.  What  that  is  we  learn  from  the  barometer. 
As  this  pressure  varies,  it  is  necessary  to  read  the  barometer 
whenever  a  gas  is  measured,  and  then  to  reduce  the  observed 
volume  to  certain  conditions  which  are  accepted  as  standard. 
If  the  gas  is  measured  in  a  tube  over  mercury  or  water,  and 
the  level  of  the  liquid  inside  the  tube  is  higher  than  that  out- 
side, the  gas  is  under  diminished  pressure,  the  amount  of 
diminution  depending  on  the  height  of  the  column  of  mer- 


MEASUREMENT  OF  THE  VOLUME  OF  OASES.       743 

cury  or  water  in  the  tube.  Thus,  if  the  arrangement  is  as 

represented  in  Fig.  18,  the  height 
of  the  mercury  column  above  the 
level  of  the  mercury  in  the  trougli  be- 
ing 100  millimeters,  and  the  pressure 
of  the  atmosphere  760  millimeters 
of  mercury;  then  the  gas  in  the  tube 
is  plainly  not  under  the  full  atmos- 
pheric pressure,  for  the  atmosphere 
is  supporting  a  column  of  mercury 
100  millimeters  high,  and  the  pres- 
sure actually  brought  to  bear  on  the 
gas  corresponds  to  760  —  100  =  660 
mm.  of  mercury.  Suppose  that  in 
this  case  the  volume  of  gas  actually 
FIG.  is.  measured  is  75  cc.  Call  this  V. 

What  would  be  the  actual  volume  V  under  the  standard  760 

mm.  ?    We  have  seen  that 


V'P'. 
Now,  in  this  case  P  =  760,  V  =  75,  and  P  '  =  660,    There- 


fore, 760  V  =  75  X  660,  or  V=  Q        =  65.13. 

In  all  cases  it  is  necessary  to  make  a  correction  similar  to 
this  in  dealing  with  the  volumes  of  gases.  The  correction  for 
temperature  and  that  for  pressure  may  be  made  in  one  opera- 
tion, the  formula  being 

273  V'P'  V'P' 

T/   —    _____  _  f\-r*        T/    — 

'  ' 


760(273  +  *)'  '   760(1  +  0.00366*)' 

in  which  V  =  the  volume  of  the  gas  at  0°  and  760  mm.  pres- 
sure ;  V  =  the  observed  volume  ;  t  =  the  observed  temper- 
ature ;  P'  =  the  pressure  under  which  the  .gas  is  measured. 
Some  of  the  most  important  ideas  which  have  been  intro- 
duced into  chemistry  with  a  view  of  explaining  the  regularities 
observed  in  the  quantities  of  substances  which  act  upon  one 
another  chemically  have  their  origin  in  observations  on  the 
conduct  of  gases.  It  is  therefore  of  the  highest  importance 
that  the  student  should  familiarize  himself  with  the  meaning 
of  the  expression,  "  the  volume  of  a  gas  under  standard  condi- 
tions." The  presence  of  water  vapor  in  a  gas  also  influences 


744      EXPERIMENTS  TO  ACCOMPANY  CHAPTER  II. 


its  volume,  and  this  must  be  taken  into  account  in  refined 
work.  The  formula  for  making  all  the  corrections  required 
in  determining  the  volume  of  a  gas  is 


FIG.  19. 


760(273  +  t)  ' 


or     V= 


V'(P'-d) 
760(1  +  0.003660' 


in  which  the  letters  F,  V,  P',  and  t 
have  the  same  significance  as  in  the  last 
formula  given,  while  a  is  the  tension  of 
water  vapor  at  t°. 

A  convenient  apparatus  for  meas- 
uring gas  volumes,  which  simplifies 
the  process,  is  that  represented  in 
Fig.  19.  It  consists  of  two  tubes  con- 
nected at  the  base  by  means  of  a  piece 
of  rubber  tubing,  and  containing  water. 
The  tube  A  is  graduated,  the  other 
is  not.  The  gas  the  volume  of  which 
is  to  be  measured  is  brought  into  the 
tube  A,  with  the  narrow  opening  at 
the  top,  and  the  other  tube  is  then 
placed  at  the  side  of  the  one  con- 
taining the  gas,  and  its  height  ad- 
justed so  that  the  column  of  liquid 
in  both  tubes  is  at  the  same  level. 
Under  these  circumstances,  obviously, 
the  gas  is  under  the  atmospheric  pres- 
sure for  which  the  necessary  correction 
must  of  course  be  made.  It  is  also 
necessary  in  this  case  to  make  the  cor- 
rections for  temperature  and  the  ten- 
sion of  aqueous  vapor.  It  is,  further, 
sometimes  convenient  when  the  gas  is 
measured  over  water  to  transfer  the 
measuring-tube  to  a  vessel  containing 
enough  water  to  permit  the  immersion 
of  the  tube  to  a  point  at  which  the  level 
of  the  liquid  inside  and  outside  of  the 
tube  is  the  same.  In  this  case  the 
conditions  are  the  same  as  in  the  apparatus  described  in  the 
last  paragraph.  The  arrangement  is  represented  in  Fig.  20. 


DECOMPOSITION  OF  POTASSIUM  CHLORATE.        745 

DETERMINATION 'OF  THE  AMOUNT  OF   OXYGEN   LIBERATED 
WHEN  A  KNOWN  WEIGHT  OF   POTASSIUM  G'HLORATE  is_ 

DECOMPOSED    BY    HEAT. 

Experiment  17.— To  determine  how  much  oxygen  is  given 
off  when  a  known  weight  of  potassium  chlorate  is  decomposed 
by  heat,  proceed  as  follows:  In  a  small  dry  hard-glass  tube 
about  10  cm.  long  and  5  to  7  mm.  internal  diameter,  closed 
at  one  end,  weigh  out  on  a  chemical  balance  about  0.25  gram 
dry  potassium  chlorate,  first  weighing  the  tube  empty.     In- 
troduce just  above  the  potassium  chlorate  a  plug  of  asbestos 
which  has  been  ignited,  then,  by  means  of  a 
blast -lamp,    soften     the    upper    end    of    the 
tube,   and   draw   it   out   so    that    it    has    the 
form  shown  in  Fig.  21.     Now  weigh  the  tube 
again. 
Let 

a  =  weight  of  tube  empty  ; 

b  =  weight  of  tube  with  potassium  chlorate  ; 

c  =  weight  of  tube  with  potassium  chlorate 

and  plug. 

Connect  at   A   by  means   of  a  short  piece  of 
rubber  tubing  with  the  measuring  tube  Fig.  19 
B 


FIG.  20.  FIG.  21. 

so  that  the  ends  of  the  two  tubes  are  almost  in  contact  with 
each  other,  the  measuring  tube  having  been  previously  filled 
with  water  to  the  zero  point,  and  the  top  closed  by  means 
of  the  stop-cock.  Open  the  stop-cock,  and  now  heat  the  po- 
tassium chlorate  gently  at  first,  and  gradually  higher  until  no 
more  gas  is  given  off.  After  the  gas  has  stood  for  half  an  hour 
to  cool  it  down  to  the  temperature  of  the  air,  adjust  the  two 
tubes  of  the  measuring  apparatus  so  that  the  level  of  the  wa- 
ter in  both  is  the  same;  read  off  the  volume  of  gas.  At  the 
same  time  read  the  barometer  and  thermometer;  and  now 
make  the  corrections  for  pressure  and  temperature  as  above 
directed.  The  weight  of  a  liter  or  1000  cc.  of  oxygen  at  0° 
and  760  mm.  pressure  is  1.4290  grams.  Knowing  the  volume 
of  oxygen  obtained,  calculate  the  weight  of  this  volume. 


746      EXPERIMENTS  TO  ACCOMPANY  CHAPTER  II. 


Remove  the  tube  containing  the  product  left  after  the  decom* 
position  of  the  potassium  chlorate,  and  weigh  it. 

Let 

d  =  weight  of  tube  after  decomposition  of  potassium  chlorate. 
Now 

b  —  a  =  weight  of  potassium  chlorate  used  ; 

d  —  (a  -\-  c  —  1}  =  weight  of  potassium  chloride  left. 

Knowing  further  the  weight  of  the  oxygen  obtained  in  the 

decomposition,  which  we  may  call  e,  it  is  obvious  from  what  has 

been  said  that 

d  —  (a  -(-  c  —  V)  -\-  e  should  be  equal  to  b  —  at 

and  the  weights  should  all  be  in  accordance  with  the  equation 

KC103  =  KC1  +  30. 
Make  all  the  calculations,  and  see  how  nearly  the  results 

obtained  agree  with  what  is  required  by  this  equation.     Should 

the  results  not  be  satisfactory  the  first  time,  repeat  the  work. 

The  more  carefully  the  work  is  done,  the  more  nearly  will  the 

results  agree  with  the  equation. 
Experiment  18. — Mix  25  to  30  grams  (or  about  an  ounce) 

of  potassium  chlorate  with  an  equal  weight  of  manganese  di- 
oxide in  a  mortar.  The  sub- 
stances need  not  be  in  the 
form  of  powder.  Heat  the 
mixture  in  a  glass  retort,  and 
collect  the  gas  by  displace- 
ment of  water  in  appropri- 
ate .  vessels, — cylinders,  bell- 
glasses,  bottles  with  wide 
mouths,  etc.  It  will  also  be 
well  to  collect  some  in  a 
gasometer,  such  as  is  com- 
monly found  in  chemical  lab- 
oratories, the  essential  features 
of  which  are  represented  in 
Fig.  22.  It  is  made  either  of 
metal  or  of  glass.  The  open- 
ing at  d  can  be  closed  by 
means  of  a  screw  cap.  In 
order  to  fill  it  with  water, 
open  the  stop-cocks  and  pour 
FIG.  22.  the  water  into  the  upper  part 

of  the  vessel  after  having  screwed  on  the  cap  d.    When  it  is 


PHYSICAL  PROPERTIES  OF  OXYGEN. 


747 


full,  water  will  flow  out  of  the  small  tube  e.  Now  close  all 
the  stop  cocks,  and  remove  the  cap  d.  The  water  will  stay 
in  the  vessel  for  the  same  reason  that  it  will  stay  in  a 
cylinder  inverted  with  its  mouth  below  water.  To  fill  the 
gasometer  with  gas,  put  it  over  a  tub  or  sink,  and  introduce 
the  tube  from  which  gas  is  issuing  into  the  opening  at  d. 
The  gas  will  rise  and  displace  the  water,  which  will  flow  out 
at  d.  When  full,  put  the  cap  on.  To  get  the  gas  out  of  the 
gasometer,  attach  a  rubber  tube  to  e,  pour  water  into  the 
upper  part  of  the  gasometer,  open  the  stop-cock  a  and  that 
at  e,  when  the  gas  will  flow  out,  and  the  current  can  be  regu- 
lated by  means  of  the  stop-cock  at  e. 

The  arrangement  of  the  retort  is  shown  in  Fig.  23. 


FIG.  23. 


FIG.  24. 


PHYSICAL  PROPERTIES  OF  OXYGEN. 

Experiment  19. — Inhale  a  little  of  the  gas  from  one  of  the 
bottles.     Has  it  any  taste  ?  any  odor  ?  any  color  ? 


CHEMICAL  PROPERTIES  OF  OXYGEN. 

Experiment  2O. — Turn  three  of  the  bottles  containing  oxy- 
gen with  the  mouth  upward,  leaving  them  covered  with  glass 
plates.  Into  one  introduce  some  sulphur  in  a  so-called  de- 
flagrating-spoon,  which  is  a  small  cup  of  iron  or  brass  at- 
tached to  a  stout  wire  which  passes  through  a  metal  plate, 
usually  of  tin  (see  Fig.  24).  In  another  put  a  little  char- 


748      EXPERIMENTS  TO  ACCOMPANY  CHAPTER  II. 

coal  (carbon),  and  in  a  third  a  piece  of  phosphorus*  about 
the  size  of  a  pea.  Let  them  stand  quietly,  and  notice  what 
changes,  if  any,  take  place.  Sulphur,  carbon,  and  phos- 
phorus are  elements,  and  oxygen  is  an  element.  It  will  be 
noticed  that  the  sulphur  and  the  carbon  remain  unchanged, 
while  some  change  is  taking  place  in  the  vessel  containing  the 
phosphorus,  as  is  shown  by  the  appearance  of  white  fumes. 
After  some  time  the  phosphorus  will  disappear  entirely,  the 
fumes  will  also  disappear,  and  there  will  be  nothing  to  show 
us  what  has  become  of  the  phosphorus.  If  the  temperature 
of  the  room  is  rather  high,  it  may  happen  that  the  phosphorus 
takes  fire.  If  it  should,  it  will  burn  with  an  intensely  bright 
light.  After  the  burning  has  stopped,  the  vessel  will  be  filled 
with  white  fumes,  but  these  will  quickly  disappear,  and  the 
vessel  will  apparently  be  empty.  What  do  these  experiments 
prove  with  reference  to  the  action  of  oxygen  on  sulphur,  car- 
bon, and  phosphorus  at  the  ordinary  temperature  ? 

Experiment  21. — In  a  deflagrating-spoon  set  fire  to  a  little 
sulphur  and  let  it  burn  in  the  air.  Notice  whether  it  burns 
with  ease  or  with  difficulty.  Notice  the  odor  of  the  fumes 
which  are  given  off.  Now  set  fire  to  another  small  portion, 
and  introduce  it  in  a  spoon  into  one  of  the  vessels  containing 
oxygen.  It  will  be  seen  that  the  sulphur  burns  much  more 
readily  in  the  oxygen  than  in  the  air.  Notice  the  odor  of  the 
fumes  given  off.  Is  it  the  same  as  that  noticed  when  the 
burning  takes  place  in  the  air  ? 

Experiment  22. — Perform  similar  experiments  with  char- 
coal. 

Experiment  23. — Burn  a  piece  of  phosphorus  not  larger 
than  a  pea  in  the  air  and  in  oxygen.  In  the  latter  case  the 
light  emitted  from  the  burning  phosphorus  is  so  intense 
fchat  it  is  painful  to  some  eyes.  It  is  better  to  be  cautious. 
The  phenomenon  is  an  extremely  brilliant  one.  The  walls 
of  the  vessel  in  which  the  burning  takes  place  become  cov- 


*  Phosphorus  should  be  handled  with  great  care.  It  is  always  kept 
under  water,  usually  in  the  form  of  sticks.  If  a  small  piece  is  wanted, 
take  out  a  stick  with  a  pair  of  forceps,  and  put  it  under  water  in  an 
evaporating-dish.  While  it  is  under  the  water,  cut  off  a  piece  of  the 
size  wanted.  Take  this  out  by  means  of  a  pair  of  forceps,  lay  it  for 
a  moment  on  a  piece  of  filter-paper,  which  will  absorb  most  of  the 
water;  then  quickly  put  it  in  the  spoon. 


OXYGEN  IS   USED    UP  IN  COMBUSTION.  749 

ered  with  a  white  substance,  which  afterwards  gradually  dis- 
appears. 

What  differences  do  you  notice  between  the  burning  in  the 
air  and  in  oxygen  ?  In  the  experiments  is  there  any  sulphur, 
or  carbon,  or  phosphorus  left  behind  ?  Do  the  experiments 
furnish  any  evidence  that  oxygen  takes  part  in  the  action  ?  or 
that  oxygen  is  used  up  ? 

Experiment  24. — Straighten  a  steel  watch-spring,*  and 
fasten  it  in  a  piece  of  metal — such  as  is  used  for  fixing  a 
deflagrating-spoon  in  an  upright  position;  wind  a  little  thread 
around  the  lower  end,  and  dip  it  in  melted  sulphur.  Set  fire 
to  this,  and  insert  it  into  a  vessel  containing  oxygen.  For  a 
moment  the  sulphur  will  burn  as  in  Experiment  21;  but  soon 
the  steel  begins  to  burn  brilliantly,  and  the  burning  continues 
as  long  as  there  is  oxygen  left  in  the  vessel.  Notice  that  in 
this  case  there  is  no  flame,  but  instead  very  hot  particles  are 
given  off  from  the  burning  iron.  The  phenomenon  is  of  great 
beauty,  especially  if  observed  in  a  dark  room.  The  walls  of  the 
vessel  become  covered  with  a  dark  reddish-brown  substance, 
some  of  which  will  also  be  found  at  the  bottom  in  larger 
pieces. 

OXYGEN  is  USED  UP  IN  COMBUSTION. 

Experiment  25. — Is  the  odor  of  the  contents  of  the  bottle  in 
which  the  sulphur  was  burned  the  same  as  before  the  experi- 
ment ?  Introduce  a  stick  with  a  small  flame  on  it  successively 
into  the  vessels  used  in  burning  sulphur,  carbon,  phosphorus, 
and  iron.  Is  oxygen  present  or  not  ?  What  evidence  have 
you  on  this  point  ? 

Experiment  26. — Fill  a  tube  say  30  to  40  cm.  (12  to  15 
inches)  long,  and  2-J-  to  3  cm.  (1  to  1£  inches)  wide,  with 
oxygen,  and  arrange  it  in  a  vessel  over  water,  as  shown  in 
Fig.  ^25.  Now  fasten  a  small  stick  of  phosphorus  to  the  end 
of  a  wire  and  push  it  into  the  tube  so  that  about  £  to  %  inch 
of  the  phosphorus  is  above  the  water  and  exposed  to  the 
•oxygen.  At  first  no  action  will  take  place,  but  after  a 

*  Old  watch-springs  can  generally  be  had  of  any  watch  maker  or 
mender  for  the  asking.  A  spring  can  be  straightened  by  unrolling  it, 
attaching  a  weight,  and  suspending  the  weight  by  the  spring.  The 
spring  is  then  heated  up  and  down  to  redness  with  the  flame  of  a  Bun- 
sen  burner. 


750       EXPERIMENTS  TO  ACCOMPANY  CHAPTER  II. 


time  white  fumes  will  be  seen   to  rise  from  the  phosphorus,  * 

and  the  phosphorus  will  begin  to  melt.     This 

action  will  be  accompanied  by  a  diminution 

of  the  volume  of  the  oxygen,  as  will  be  shown 

by  the  rise  of  the  water.    When  the  water  has 

risen  so  as  to  cover  the  phosphorus,  shove  the 

stick  up  so  that   it   is  again  just  above  the 

surface  of  the  water.     Some  of   the  oxygen 

will  again  be  used  up.     By  working  carefully, 

and  repeating  this  process  as  many  times  as 

may  be  necessary,  the  oxygen  can  all  be  used 

up  without  the  active  burning  of  the  phos- 

phorus.    Usually,  however,  before  the  action 

is  completed,  the  temperature  of  the  phos- 

phorus becomes  so  high  that   it  takes   fire, 

when  there  is  a  flash  of  light  in  the  tube  and 

a  sudden  rise  of  the  water,  showing  that  the 

gas  is  suddenly  used  up. 

Experiment  27.  —  Burn  a  steel  watch-spring  as  directed  in 
Experiment  24,  with  the  difference  that  the 
spring  is  passed  air-tight  through  a  cork  which 
is  fitted  tightly  into  the  neck  of  the  bell-jar. 
As  the  spring  burns,  the  water  will  rise  from 
the  vessel  in  which  the  bell-jar  is  standing, 
and  it  is  necessary  to  pour  water  into  this  ves- 
sel. When  the  spring  has  burned  near  to- 
the  cork  shove  it  through  so  that  the  burning 
may  continue.  If  the  experiment  is  properly 
performed  the  bell-  jar  will  be  nearly  full  of 
water  at  the  end.  What  does  this  prove  ? 

THE  PRODUCTS  OF  COMBUSTION  WEIGH  MORE 
THAN  THE  BODY  BURNED. 


FIG.  25. 


Flo 


Experiment  28. — Weigh  off  about  a  gram 
of  magnesium  ribbon  in  a  porcelain  crucible. 
Heat  over  a  Bunsen  burner  until  the  magnesium  has  turned 
to  a  white  substance  (magnesium  oxide).  After  cooling,  weigh 
again.  Perform  the  same  experiment  with  zinc,  tin,  and  lead. 
What  conclusion  are  you  justified  in  drawing? 

Experiment  29. — Over   each   pan   of  a   large  and  rather 
sensitive  balance  suspend  a  glass  tube  filled  with  pieces  of  solid 


PREPARATION  OF  HYDROGEN.  751 

caustic  soda.  A  balance  that  will  answer  the  purpose  very 
well  can  be  made  of  wood  with  metal  bearings.  It  may  con- 
veniently be  about  2?  feet  high,  with  a  delicate  beam  about 
3  feet  long.  The  best  tubes  for  the  caustic  soda  are  Argand 
lamp-chimneys,  around  the  bottom  of  which  is  tied  a  piece  of 
wire-gauze  to  prevent  the  caustic  soda  from  falling  out.  On 
one  pan  of  the  balance  place  a  candle  directly  under  one  of 
the  caustic-soda  tubes,  so  adjusted  that  the  flame  shall  be  not 
more  than  ty  to  3  inches  below  the  bottom  of  the  tube.  By 
means  of  weights  placed  on  the  other  pan  establish  equilibrium. 
Now  light  the  candle.  Slowly,  as  it  burns,  the  pan  upon 
which  it  is  placed  will  sink,  showing  that  the  products  of  com- 
bustion which  are  partly  absorbed  by  the  caustic  soda  are 
heavier  than  the  candle  was.  While  this  is  by  no  means  an 
accurate  experiment,  it  is  a  very  striking  one,  and  proves  be- 
yond question  that  in  the  process  of  combustion  matter  is 
taken  up  by  the  burning  body. 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  III. 

PREPARATION  OF  HYDROGEN. 
Experiment  3O. — Kepeat  Experiment  3  and  examine  the 


•  Experiment  31.— Throw  a  small  piece  of  sodium  *  on  water. 
While  it  is  floating  on  the  surface  apply  a  lighted  match  to  it. 
A  yellow  flame  will  appear.  This  is  burning  hydrogen,  the 
flame  being  colored  yellow  by  the  presence  of  the  sodium,  some 
of  which  also  burns.  Make  the  same  experiment  with  potas- 
sium. The  flame  appears  in  this  case  without  the  aid  of  the 
match.  It  has  a  violet  color,  which  is  due  to  the  burning  of 
some  of  the  potassium.  The  gas  given  off  in  these  experi- 
ments is  either  burned  at  once  or  escapes  into  the  air.  In  the 
case  of  the  potassium  it  takes  fire  at  once,  because  the  action 
takes  place  rapidly  and  the  heat  evolved  is  sufficient  to  set  fire  to 
it;  in  the  case  of  the  sodium,  however,  the  action  takes  place 
more  slowly,  and  the  temperature  does  not  get  high  enough  to 
set  fire  to  the  gas.  In  order  to  collect  it  unburned,  it  is  only 
necessary  to  allow  the  decomposition  to  take  place  so  that  the 

*  The  metals  sodium  and  potassium  are  kept  under  oil.  When  a 
small  piece  is  wanted  take  out  .one  of  the  larger  pieces  from  the  bottle, 
roughly  wipe  off  the  oil  with  filter-paper,  and  cut  off  a  piece  the  size 
needed.  It  is  not  advisable  to  use  a  piece  larger  than  a  small  pea. 


752      EXPERIMENTS  TO  ACCOMPANY  CHAPTER  III. 

gas  will  rise  in  an  inverted  vessel  filled  with  water.  For  this 
purpose  fill  a  good-sized  test-tube  with  water  and  invert  it  in 
a  vessel  of  water.  Cut  off  a  piece  of  sodium  not  larger  than  a 
pea,  wrap  it  in  a  layer  or  two  of  filter-paper,  and  with  the 
fingers  or  a  pair  of  curved  forceps  bring  it  quickly  below  the 
mouth  of  the  test-tube  and  let  go  of  it.  It  will  rise  to  the 
top,  the  decomposition  of  the  water  will  take  place  quietly, 
and  the  gas  formed,  being  unable  to  escape,  will  remain  in  the 
tube.  By  repeating  this  operation  in  the  same  tube  a  second 
portion  of  gas  can  be  made,  and  so  on  until  the  vessel  is  full. 

Examine  the  gas  and  see  whether  it  acts  like  the  hydrogen 
obtained  from  water  by  means  of  the  electric  current.  What 
evidence  have  you  that  they  are  the  same  ?  Is  this  evidence 
sufficient  to  prove  the  identity  of  the  two  ? 

The  metals  sodium  and  potassium  disappear  in  these  experi- 
ments, and  we  get  hydrogen.  What  becomes  of  the  metals  ? 
and  what  is  the  source  of  the  hydrogen  ?  If  after  the  action 
has  stopped  the  water  is  examined,  it  will  be  found  to  contain 
something  in  solution.  It  now  has  a  peculiar  taste,  which  we 
call  alkaline;  it  feels  slightly  soapy  to  the  touch;  it  changes 
certain  vegetable  colors.  If  the  water  is  evaporated  off,  a 
white  substance  remains  behind,  which  is  plainly  neither 
sodium  nor  potassium.  In  solid  form  or  in  very  concentrated 
solution  it  acts  very  strongly  on  animal  and  vegetable  sub- 
stances, disintegrating  many  of  them.  On  account  of  this 
action  it  is  known  as  caustic  soda,  or,  in  the  case  of  potas- 
sium, as  caustic  potassa. 


FIG.  27. 


Experiment  32. — Certain  metals  which  do  not  decompose 
water  at  ordinary  temperatures,  or  which  decompose  it  slowly, 
decompose  it  easily  at  elevated  temperatures.  This  is  true  of 


PREPARATION  OF  HYDROGEN. 


"53 


iron.  If  steam  is  passed  through  a  tube  containing  pieces  of 
iron  heated  to  redness,  decomposition  of  the  water  takes  place, 
and  the  oxygen  is  retained  by  the  iron,  which  enters  into  com- 
bination with  it,  while  the  hydrogen  is  liberated.  In  this  ex- 
periment a  porcelain  tube  with  an  internal  diameter  of  from 
20  to  25  mm.  (about  an  inch)  and  a  gas  furnace  are  desirable, 
though  a  hard-glass  tube  and  a  charcoal  furnace  will  answer. 
The  arrangement  of  the  apparatus  is  shown  in  Fig.  27. 

Experiment  33.— In  a  cylinder  or  test-tube  put  some  small 
pieces  of  zinc,  and  pour  upon  it  some  ordinary  hydrochloric 
acid.  If  the  action  is  brisk,  after  it  has  continued  for  a  min- 
ute or  two  apply  a  lighted  match  to  the  mouth  of  the  vessel. 
The  gas  will  take  fire  and  burn.  If  sulphuric  acid  diluted 
with  five  or  six  times  its  volume  of  water*  is  used  instead  of 
hydrochloric  acid,  the  same  result  will  be  reached.  The  gas 
evolved  is  hydrogen.  For  the  purpose  of  collecting  the  gas 
the  operation  is  best  performed  in  a  wide-mouthed  bottle,  in 


FIG.  28. 


FIG.  29. 


which  is  fitted  a  cork  with  two  holes  (see  .Fig.  28),  or  in  a 
bottle  with  two  necks  called  a  Wolffs  flask  (see  Fig.  29). 
Through  one  of  the  holes  a  funnel-tube  passes,  and  through 
the  other  a  glass  tube  bent  in  a  convenient  form. 

*  If  it  is  desired  to  dilute  ordinary  concentrated  sulphuric  acid  with 
water,  the  acid  should  be  poured  slowly  into  the  water  while  the  mix- 
ture is  constantly  stirred.  If  the  water  is  poured  into  the  acid,  the  heat 
evolved  at  the  places  where  the  two  come  in  contact  may  be  so  great  as 
to  convert  the  water  into  steam  and  cause  the  strong  acid  to  spatter. 


754      EXPERIMENTS  TO  ACCOMPANY  CHAPTER  IIL 

The  zinc  used  is  granulated.  It  is  prepared  by  melting  it 
in  a  ladle,  and  pouring  the  molten  metal  from  an  elevation 
of  four  or  five  feet  into  water.  The  advantage  of  this  form 
is  that  it  presents  a  large  surface  to  the  action  of  the  acids. 
A  handful  of  this  zinc  is  introduced  into  the  bottle,  and 
enough  of  a  cooled  mixture  of  sulphuric  acid  and  water  (1 
volume  concentrated  acid  to  6  volumes  water)  poured  upon  it 
to  cover  it.  Usually  a  brisk  evolution  of  gas  takes  place  at 
once.  Wait  for  two  or  three  minutes,  and  then  collect  some 
of  the  gas  by  displacement  of  water.  When  the  action  be- 
comes slow,  add  more  of  the  dilute  acid.  It  will  be  well  to 
fill  several  cylinders  and  bottles  with  the  gas,  and  also  a  gaso- 
meter, from  which  it  can  be  taken  as  it  is  needed  for  experi- 
ments. 

SOMETHING  BESIDES  HYDROGEN'  is  FORMED. 

Experiment  34. — After  the  action  is  over  pour  the  contents 
of  the  flask  through  a  filter  into  an  evaporating-dish,  and 
boil  off  the  greater  part  of  the  water,  so  that,  on  cooling,  the 
substance  contained  in  solution  will  be  deposited.  If  the  op- 
eration is  carried  on  properly,  the  substance  will  be  deposited 
in  regular  forms  called  crystals.  It  is  zinc  sulphate,  ZnS04, 
formed  by  the  replacement  of  the  hydrogen  of  the  sulphuric 
acid  by  zinc. 

PROBLEMS.— How  much  zinc  would  it  take  to  give  200  liters  of  hy- 
drogen ?  How  much  zinc  sulphate  would  be  formed  ?  How  much  hy- 
drogen would  be  formed  by  the  action  of  50  grams  of  zinc  on  sulphuric 
acid  ?  How  much  sulphuric  acid  would  be  used  up  ? 

DETERMINATION  OF  THE  AMOUNT  OF  HYDROGEN  EVOLVED 
WHEN  A  KNOWN  WEIGHT  OF  ZINC  is  DISSOLVED  IN  SUL- 
PHURIC ACID. 

Experiment  35. — This  determination  can  be  made  by  means 
of  an  apparatus  such  as  represented  in  Fig.  30.  The  bent 
tube  leading  from  the  flask  A  is  drawn  out  at  B,  and  a  plug 
of  glass-wool  introduced  below  the  constriction.  The  other 
parts  of  the  apparatus  need  no  description.  The  flask  should 
have  a  capacity  of  about  40  to  50  cc. ;  and  the  measuring  tube 
C should  have  a  capacity  of  about  100  cc.,  and  be  graduated  to 
•  cc. 


AMOUNT  OF  HYDROGEN  EVOLVED. 


'55 


"  The  experiment  is  conducted  in  the  following  manner  : 
D  is  filled  with  distilled  water  ;  a  piece  of  zinc  weighing  from 
0.150  to  0.200  gram  is  placed  in  the  flask;  the  pinch-cock  E 
is  then  opened,  and  the  whole  apparatus  thus  filled  with 
water.  The  apparatus  is  now  examined  in  order  to  ascertain 
if  gas  bubbles  are  lodged  under  the  stopper  F  or  in  the  glass- 
wool.  If  so,  they  can  usually  be  dislodged  without  difficulty. 
If  they  persist,  a  few  moments'  boiling  of  the  water  in  the 
flask  will  eifect  their  complete  removal.  .  .  The  eudiometer  is 
now  placed  over  the  outlet  of  the  delivery-tube,  and  the 
greater  portion  of  the  water  remaining  in  D  allowed  to  flow 
through  the  apparatus.  Sulphuric  acid  of  the  concentration 
ordinarily  employed  in  the  laboratory  (1  of  H2S04  to  4  of  H30) 
is  poured  into  the  reservoir  D  until  it  is  nearly  full.  The 
pinch-cock  E  is  then  opened,  and  the  water  which  fills  the 


Fio.  30. 

apparatus  is  displaced  by  sulphuric  acid.  The  action  of  the 
acid  upon  the  metal  may  be  facilitated  by  heat  or  by  adding 
some  platinum  scraps.  When  the  action  is  over,  the  contents 
of  the  flask  are  swept  through  the  delivery-tube  by  again  open- 
ing the  pinch-cock  E.  Finally,  the  measuring-tube  is  trans- 
ferred to  a  cylinder  of  water,  and  the  volume  of  the  gas  read 
and  corrected  in  the  usual  manner.  If  hydrochloric  instead  of 
sulphuric  acid  has  been  used,  which  would  be  the  case  when 
the  metal  employed  is  aluminium,  a  little  caustic  soda  should 
be  added  to  the  water  in  the  cylinder  to  which  the  eudiometer 
is  transferred."*  • 


*  See  Morse  and  Keiser,  American  Chemical  Journal,  vol.  vi.  p.  349. 


?5ti     EXPERIMENTS  TO  ACCOMPANY  CHAPTER  III. 

A  liter  of  hydrogen  at  0°  and  760  mm.  weighs  0.089873 
gram.  How  much  does  the  hydrogen  obtained  in  the  experi- 
ment weigh?  How  much  ought  to  have  been  obtained? 
How  many  cubic  centimeters  of  hydrogen  ought  to  have  been 
obtained  ? 

Try  the  same  experiment,  using  tin-  and  hydrochloric  acid. 
The  action  takes  place  as  represented  in  the  equation 

Sn  +  2HC1  =  SnCla  +  H2. 

It  would  be  well,  further,  to  try  the  experiment  also  with  iron 
and  sulphuric  acid,  and  with  aluminium  and  hydrochloric 
acid,  and  to  calculate  from  the  results  the  relation  between 
the  weights  of  the  four  metals  required  to  give  equal  volumes 
of  hydrogen,  and  the  volumes  of  hydrogen  given  by,  say,  a 
gram  of  each  metal.  The  action  between  iron  and  sulphuric 
acid  takes  place  according  to  the  equation 

Fe  +  H2S04  =  FeS04  +  Ha. 

That  between  aluminium  and  hydrochloric  acid  is  represented 
by  this  equation  : 

Al  +  3HC1  =  A1C13  +  3H. 

HYDROGEN  is  PURIFIED  BY  PASSING  THROUGH  A  SOLUTION 

OF  POTASSIUM  PERMANGANATE. 

Experiment  36. — Pass  some  of  the  gas,  made  by  the  action 
of  zinc  on  sulphuric  acid,  through  a  wash  cylinder  contain- 


FIG.  31. 

ing  a  solution  of  potassium  permanganate;  collect  some  of  it, 
and  notice  whether  it  has  an  odor.     The  apparatus  should 


DIFFUSION. 


75? 


Fia.  32. 


be  arranged  as  shown  in  Fig.  31.  The  solution  of  potas- 
sium permanganate  is,  of  course,  contained  in  the  small  cyl- 
inder A,  and  the  tubes  so  arranged  that  the  gas  bubbles 
through  it. 

Has  the  gas  any  odor  or  taste  or  color  ? 

Experiment  37. — Place  a  vessel  containing  hydrogen  with 
the  mouth  upward  and  uncovered.  In  a  short  time  examine 
the  gas  contained  in  the  vessel,  and  see  whether  it  is  hydrogen, 
What  does  this  experiment  prove  with  reference  to  the  weight 
of  hydrogen  as  compared  with  that  of  the  air  ? 

Experiment  38. — Gradually  bring  a  vessel  containing  hy- 
drogen with  its  mouth  upward 
below  an  inverted  vessel  contain- 
ing air,  in  the  way  shown  in  Fig. 
32.  After  the  vessel  which  con- 
tained the  hydrogen  has  been 
brought  in  the  upright  position 
beneath  the  other,  examine  the 
gas  in  each  vessel.  Which  one 
contains  the  hydrogen  ? 

Experiment  39. — Soap-bubbles  filled  with  hydrogen  rise  in 
the  air.  This  experiment  is  best  performed  by  connecting  an 
ordinary  clay  pipe  by  means  of  a  piece  of  rubber  tubing  with 
the  delivery-tube  of  a  gasometer  filled  with  hydrogen.  Small 
balloons  of  collodion  are  also  made  for  the  purpose  of  showing 
the  lightness  of  hydrogen. 

HYDROGEN  PASSES  READILY  THROUGH  POROUS  VESSELS. 
DIFFUSION. 

9 

Experiment  4O. — Arrange  an  apparatus  as  shown  in  Fig. 
S3.  It  consists  of  a  porous  earthenware  cup,  such  as  is  used 
in  galvanic  batteries,  fitted  wi'th  a  perforated  cork  connected 
with  a  glass  tube  2  to  3  feet  long.  The  cork  must  fit  air-tight 
into  the  mouth  of  the  cup,  as  well  as  the  tube  into  the  cork0 
This  may  be  secured  by  shoving  the  cork  into  the  cup  until 
its  outer  surface  is  even  with  the  edge  of  the  cup,  and  then 
covering  it  carefully  with  sealing-wax.  Put  the  lower  end  of 
the  glass  tube  through  a  cork  into  one  neck  of  a  Wolffs 
bottle  containing  some  water  colored  with  litmus  or  indigo,  so 
that  the  end  of  the  tube  is  above  the  surface  of  the  water. 
Through  the  other  neck  of  the  bottle  pass  a  tube  slightly  bent 
outward  and  drawn  out  at  the  end  to  a  fine  opening.  This 


758      EXPERIMENTS  TO  ACCOMPANY  CHAPTER  III. 


tube  must  also  be  fitted  to  the  bottle  by  an  air- tight  cork, 

and  its  lower  end  must  be 
below  the  surface  of  the 
liquid.  Now  bring  a  bell- 
jar  containing  dry  hydro- 
gen over  the  porous  cup, 
when  the  liquid  will  be 
seen  to  rise  in  the  short, 
bent  tube  that  dips  be- 
low the  liquid,  and  be 
forced  out  of  it,  some- 
times with  considerable 
velocity.  Withdraw  the 
bell-jar,  and  bubbles  will 
rise  rapidly  from  the  bot- 
tom of  the  tube  which 
dips  under  the  water,  thus 
showing  that  air  is  enter- 


FIG.  33. 


FIG.  34. 


ing  the  bottle.  This  is  due  to  the  diffusion  of  the  hydrogen 
from  the  porous  cup  into  the  air.  Explain  all  that  you  have 
seen. 

CHEMICAL  PROPERTIES  OF  HYDROGEN. 

Experiment  41. — If  there  is  no  small  platinum  tube  avail- 
able, roll  up  a  small  piece  of  platinum-foil  and  melt  it  into 
the  end  of  a  glass  tube,  as  shown  in  Fig.  34.  Connect  the 
burner  thus  made  with  the  gasometer  containing  hydrogen, 
and  after  the  gas  has  been  allowed  to  issue  from  it  for  a 
moment,  set  fire  to  it.  In  a  short  time  it  will  be  seen  that  the 
flame  is  practically  colorless,  and  gives  no  light.  That  it  is 


PRODUCT  FORMED   WHEN  HYDROGEN  IS  BURNED.    759 

hot  can  be  readily  shown  by  holding  a  piece  of  platinum  wire 
or  a  piece  of  some  other  metal  in  it. 

Experiment  42. — Into  the  flame  of  burning  hydrogen  in- 
troduce a  small  coil  of  platinum  wire.  What  change  is  ob- 
served ?  Introduce  also  a  piece  of  magnesium  ribbon.  Explain 
the  difference  between  the  two  cases.  What 
becomes  of  the  magnesium  ?  of  the  platinum  ? 

Experiment  43. — Hold  a  cylinder  filled 
with  hydrogen  with  the  mouth  downward. 
Insert  into  it  a  lighted  taper  held  on  a  bent 
wire,  as  shown  in  Fig.  35.  The  gas  takes  fire 
at  the  mouth  of  the  vessel,  but  the  taper  is 
extinguished.  On  withdrawing  the  taper  and 
holding  the  wick  for  a^ moment  in  the  burn- 
ing hydrogen,  it  will  take  fire,  but  on  putting 
it  back  in  the  hydrogen  it  will  again  be  extin- 
guished. Other  burning  substances  should 
be  tried  in  a  similar  way.  What  conclu- 
sions are  justified  by  the  last  two  experi- 
ments ?  FIG.  3t>. 

PRODUCT  FORMED  WHEN  HYDROGEN  is  BURNED. 

Experiment  44. — Hold  a  clean,  dry  glass  plate  a  few  inches 
above  a  hydrogen  flame.  What  do  you  observe?  Kemove 
what  is  deposited  upon  the  plate,  and  hold  the  plate  again 
over  the  flame.  Repeat  this  a  number  of  times.  What  does 
the  substance  deposited  upon  the  plate  suggest?  Can  you 
positively  say  what  it  is  ? 

REDUCTION. 

Experiment  45. — Arrange  an  apparatus  as  shown  in  Fig. 
36.  The  flask  A  contains  zinc  and  dilute  sulphuric  acid;  the 
cylinder  B  a  solution  of  potassium  permanganate;  the  cylinder 
0  concentrated  sulphuric  acid;  and  the  tube  D  granulated 
calcium  chloride.  The  object  of  the  potassium  permanganate 
is  to  purify  the  hydrogen;  the  object  of  the  concentrated  sul- 
phuric acid  and  calcium  chloride  is  to  remove  moisture  from 
the  gas.  In  the  tube  E  put  a  few  pieces  of  the  black  oxide  of 
copper,  or  cupric  oxide,  CuO.  After  hydrogen  has  been  pass- 
ing long  enough  to  drive  all  the  air  out  of  the  apparatus 
(about  two  or  three  minutes  if  there  is  a  brisk  evolution)  heat 


760      EXPERIMENTS  TO  ACCOMPANY  CHAPTER  IV. 

the  oxide  of  copper  by  means  of  a  flame  applied  to  the  tube. 
What  change  in  color  takes  place  ?    Try  the  action  of  nitric 


Fio.  36. 

acid  on  the  substance  before  the  action  and  after,  and  note 
whether  there  is  any  difference.  What  appears  in  G  ?  Ex- 
plain what  you  have  seen. 

Experiment  46. — Try  the  experiment  just  described,  using 
ferric  oxide,  or  oxide  of  iron,  Fe203,  instead  of  cupric  oxide. 
What  is  the  common  feature  in  the  two  reactions  ? 


EXPERIMENTS  TO  ACCOMPANY   CHAPTER  IV. 
COMPOSITION  OF  WATER. 

Experiment  47. — Arrange  the  apparatus  shown  in  Fig.  36 
with  a  straight  tube  instead  of  the  bent  tube  E,  and  connect 
this  with  a  small  bent  tube  containing  calcium  chloride,  as 
shown  in  Fig.  37.  Weigh  tube  E  empty,  and  after  the  cupric 


u 


FIG.  37. 


oxide  has  been  put  into  it.  This  gives  the  weight  of  the  cupric 
oxide.  Weigh  the  tube  F  before  the  experiment.  Now  pro- 
ceed as  in  Experiment  45.  In  this  case  all  the  water  formed 
by  the  action  of  the  hydrogen  on  the  cupric  oxide  will  be 


COMPOSITION  OF  WATER.  761 

absorbed  by  the  calcium  chloride  in  tube  F.  This  tube  will 
therefore  gain  in  weight,  and  as  oxygen  is  removed  from  the 
cupric  oxide,  tube  E  will  lose  in  weight.  After  the  reduction 
is  complete  weigh  tube  E  and  tube  F  again. 

Let  x  —  weight  of  tube  E  -f-  cupric  oxide  before  the  ex- 
periment; 

y  =  weight  of  tube  E  -{•  copper  after  the  experiment. 
Then  x  —  y  =  weight  of  oxygen   removed  from  the  cupric 

oxide. 

Let  a  =  weight  of  tube  F  before  the  experiment, 
and      b  =       "  "     "  after      "  " 

Then   b  —  a  =  weight  of  water  formed. 

If  the  experiment  is  properly  performed,  it  will  be  found 
that  the  ratio  -r — -  is  very  nearly  -.  Or  the  result  may  be 

o  —  ci>  y 

stated  thus:  In  nine  parts  of  water  there  are  eight  parts  of 
oxygen. 

Experiment  48. — The  tubes  in  the  apparatus  used  in  Ex- 
periment 3,  or  some  other  similar  apparatus,  should  be 
graduated.  Let  the  gases  formed  by  the  action  of  the  electric 
current,  as  in  Experiment  3,  rise  in  the  tubes,  and  observe 
the  volumes.  It  will  be  seen  that  when  one  tube  is  just  full 
of  gas,  the  other,  if  it  is  of  the  same  size,  will  be  only  half  full. 
On  examining  the  gases  the  larger  volume  will  be  found  to  be 
hydrogen,  and  the  smaller  volume  oxygen.  What  are  the 
relative  weights  of  equal  volumes  of  hydrogen  and  oxygen  ? 
In  what  proportion  by  weight  are  the  two  gases  obtained 
from  water  in  this  experiment?  How  does  this  result  agree 
with  that  obtained  in  the  preceding  experiment  ?  Does  this 
experiment  prove  that  water  consists  only  of  hydrogen  and 
oxygen  ? 

Experiment  49. — Pass  hydrogen  from  a  generating- flask  or 
a  gasometer  through  a  tube  containing  some  substance  that 
will  absorb  moisture  ;  for  all  gases  made  in  the  ordinary 
way  and  collected  over  water  are  charged  with  moisture. 
The  calcium  chloride  should  be  in  granulated  form,  not 
powdered.  After  passing  the  hydrogen  through  the  cal- 
cium chloride,  pass  it  through  a  tube  ending  in  a  narrow 
opening,  and  set  fire  to  it.  If  now  a  dry  vessel  is  held  over 
the  flame,  drops  of  water  will  condense  on  its  surface  and  run 


762       EXPERIMENTS  TO  ACCOMPANY  CHAPTER  IV. 

down.    A  convenient  arrangement  of  the  apparatus  is  shown 
in  Pig.  38. 


FIG.  38. 

A  is  the  calcium  chloride  tube.  Before  lighting  the  jet, 
hold  a  glass  plate  in  the  escaping  gas,  and  see  whether  water 
is  deposited  on  it.  Light  the  jet  before  putting  it  under  the 
bell- jar ;  otherwise,  if  hydrogen  is  allowed  to  escape  into  the 
vessel,  it  will  contain  a  mixture  of  air  and  hydrogen,  and  this 
mixture,  as  we  shall  soon  see,  is  explosive. 

Experiment  50. — Mix  hydrogen  and  oxygen  in  the  propor- 
tions of  about  2  volumes  of  hydrogen  to  1  volume  of  oxygen, 
in  a  gasometer.  Fill  soap-bubbles,  made  as  directed  in  Ex- 
periment 39,  with  this  mixture,  and  allow  them  to  rise  in  the 
air.  As  they  rise,  bring  a  lighted  taper  in  contact  with  them, 
when  a  sharp  explosion  will  occur.  Great  care  must  be  taken 
to  keep  all  flames  away  from  the  vicinity  of  the  gasometer 
while  the  mixture  is  in  it.  This  experiment  is  conveniently 
performed  by  hanging  up,  about  six  to  eight  feet  above  the 
experiment-table,  a  good-sized  tin  funnel-shaped  vessel,  with 
the  mouth  downward.  a  Now  place  a  gas  jet  or  a  small  flame 
of  any  kind  at  the  mouth  of  the  vessel.  If  the  soap-bubbles 
are  allowed  to  rise  below  this  apparatus  they  will  come  in  con- 
tact with  the  flame  and  explode  at  once.*  What  does  this 
experiment  show  ?  Does  it  give  any  information  in  regard  to 
the  composition  of  water  ? 

*  The  same  apparatus  may  be  used  in  experimenting  with  soap 
bubbles  filled  with  hydrogen. 


EUDIOMETRIC  EXPERIMENTS.  763 


EUDIOMETRIC  EXPERIMENTS. 

Experiment  51. — The  general  method  of  studying  the 
combination  of  hydrogen  and  oxygen  by  means  of  the  eudi- 
ometer was  described  in  the  text  (see  p.  50).  To  what  was 
there  said  it  need  only  be  added  that,  in  exploding  the  mix- 
ture in  the  eudiometer,  the  latter  should  be  held  down  firmly, 
by  means  of  a  clarnp,  against  a  thick  piece  of  rubber  cloth 
placed  on  the  bottom  of  the  mercury-trough.  In  making  the 
measurements  of  the  volume  of  the  gases  and  the  height  of 
the  mercury  column,  care  must  be  taken  to  have  the  eudiom- 
eter in  a  perpendicular  position.  This  can  be  secured  by 
means  of  plumb-lines  suspended  from  the  ceiling  and  reach- 
ing nearly  to  the  table,  by  which  the  position  of  the  eudiom- 
eter can  be  adjusted. 

OXYHYDROGEN   BLOW-PIPE. 

Experiment  52.— Hold  in  the  flame  of  the  oxyhydrogen 
blow-pipe  successively  a  piece  of  iron  wire,  a  piece  of  a  steel 
watch-spring,  a  piece  of  copper  wire,  a  piece  of  zinc,  a  piece 
of  platinum  wire. 

Experiment  53. — Cut  a  piece  of  lime  of  convenient  size 
and  shape,  say  an  inch  long  by  three  quarters  of  an  inch  wide, 
and  the  same  thickness.  Fix  it  in  position  so  that  the  flame 
of  the  oxyhydrogen  blow-pipe  will  play  upon  it.  The  light 
is  very  bright,  but  by  no  means  as  intense  as  the  electric  light. 

EXPERIMENTS  TO  ACCOMPANY   CHAPTER  V. 

ORGANIC  SUBSTANCES  CONTAIN  WATER. 

Experiment  54. — In  dry  test-tubes  heat  gently  various  or- 
ganic substances  as  a  piece  of  wood,  fresh  meat,  fruits,  vege- 
tables, etc. 

WATER  OF  CRYSTALLIZATION. 

Experiment  55. — Take  some  of  the  crystals  of  zinc  sul- 
phate obtained  in  Experiment  34.  Spread  them  out  on  a 
layer  of  filter-paper,  and  finally  press  two  or  three  of  them 
between  folds  of  the  paper.  Examine  them  carefully.  They 
appear  to  be  quite  dry,  and  in  the  ordinary  sense  they  are 
dry.  Heat  them  in  a  dry  tube,  when  it  will  be  observed  that 
water  condenses  in  the  upper  part  of  the  tube,  while  the 


764        EXPERIMENTS  TO  ACCOMPANY  CHAPTER   V. 

crystals  lose  their  lustre,  becoming  white  and  opaque,  and  at 
last  crumbling  to  powder. 

Experiment  56. — Perform  a  similar  experiment  with  some 
gypsum,  which  is  the  natural  substance  from  which  "  plaster 
of  Paris  "  is  made. 

Experiment  57. — Heat  a  few  small  crystals  of  copper  sul- 
phate, or  blue  vitriol.  In  this  case  the  loss  of  water  is  accom- 
panied by  a  loss  of  color.  After  all  the  water  is  driven  off, 
the  powder  left  behind  is  white.  On  dissolving  it  in  water, 
however,  the  solution  will  be  seen  to  be  blue  ;  and  if  the  solu- 
tion is  evaporated  until  the  substance  is  deposited,  it  will 
again  appear  in  the  form  of  blue  crystals. 
EFFLORESCENT  SALTS. 

Experiment  58. — Select  a  few  crystals  of  sodium  sulphate 
which  have  not  lost  their  lustre.  Put  them  on  a  watch- 
glass,  and  let  them  lie  exposed  to  the  air  for  an  hour  or  two. 
They  soon  lose  their  lustre,  and  undergo  the  changes  noticed 
in  heating  zinc  sulphate. 

DELIQUESCENT  SALTS. 

Experiment  59. — Expose  a  few  pieces  of  calcium  chloride 
to  the  air.  Its  surface  will  soon  give  evidence  of  the  presence 
of  moisture,  and  after  a  time  the  substance  will  dissolve  in 
the  water  which  is  absorbed. 

PURIFICATION   OF   WATER  BY  DISTILLATION". 

Experiment  6O. — In  an  apparatus  like  that  shown  in  Fig. 
39  distil  a  dilute  solution  of  copper  sulphate  or  some  other 


FIG.  39. 

colored  substance.     A  slow  current  of  cold  water  must  be 


METHOD  OF  DUMA&  765 

kept  running  through  the  condenser  by  connecting  the  lower 
rubber  tube  with  a  water-cock.  When  the  water  is  boiled  in 
the  large  flask,  the  steam  passes  into  the  inner  tube  of  the 
condenser.  As  this  is  surrounded  by  cold  water,  the  steam 
condenses  and  the  distilled  water  collects  in  the  receiver. 

EXPERIMENTS   TO  ACCOMPANY   CHAPTER  VI. 

It  would  be  well  in  this  connection  to  determine  the  specific 
gravity  of  some  substance  in  the  form  of  vapor.  The  princi- 
pal methods  for  this  purpose  are  those  of  Dumas,  Gay  Lussac, 
Hof  mann,  and  Victor  Meyer.  That  of  Dumas,  which  consists 
in  measuring  the  volume  and  determining  the  weight  of  the 
vapor  under  observation,  is  the  most  accurate.  The  method 
of  Hofmann  is  a  modification  of  that  of  Gay  Lussac.  It  con- 
sists in  weighing  a  small  quantity  of  the  liquid  the  specific 
gravity  of  whose  vapor  is  to  be  determined,  and,  after  intro- 
ducing the  liquid  in  a  minute  glass  vessel  into  a  eudiometer 
over  mercury,  heating  the  eudiometer  and  its  contents  by 
passing  steam  through  a  jacket  surrounding  it  and  measuring 
the  volume  of  vapor  formed.  The  method  of  Victor  Meyer 
is  used  very  commonly,  especially  when  it  is  required  to 
-determine  the  specific  gravity  of  the  vapor  of  a  substance 
which  boils  at  a  high  temperature. 

METHOD  OF  DUMAS. 

Experiment  61. — In  this  method  the  liquid  to  be  vaporized 
is  brought  into  a  small  balloon  like  that  shown  in  Fig.  40. 
The  dry  balloon  is  first  weighed,  and  a  small  quantity  of 
liquid  then  introduced  by  gently  heating  the  balloon  and  pat- 
ting the  point  of  its  stem  into  the  liquid,  when,  on  cooling,  the 
liquid  rises  and  enough  is  easily  brought  into  the  balloon  in 
this  way.  The  balloon  is  now  placed  (in  the  position  shown  in 
Pig.  41)  in  a  bath  of  water,  oil,  or  paraffin,  according  to  the 
boiling-point  of  the  liquid.  The  bath  is  heated  30-40°  above 
the  boiling-point  of  the  liquid  under  examination.  The  air 
is  thus  driven  out  and  the  balloon  is  filled  with  the  vapor. 
When  vapor  no  longer  escapes,  the  point  of  the  stem  is  closed 
by  melting  it  with  a  mouth  blow-pipe.  The  balloon  is  then 
cleaned,  dried,  and  weighed.  The  temperature  of  the  bath 
and  the  height  of  the  barometer  are  observed  at  the  time  the 
balloon  is  closed.  The  point  of  the  stem  is  broken  off  under 


766      EXPERIMENTS  TO  ACCOMPANY  CHAPTER   VI. 


mercury,  when  the  mercury  rises  and  fills  the  balloon.  By 
pouring  the  mercury  out  into  a  graduated  cylinder  the  ca- 
pacity of  the  balloon  is  determined.  The  specific  gravity  of 
the  vapor  is  calculated  by  the  aid  of  the  formula 


B  +  p)(l  +  0.00366  X 


vhv  X  0.001293 


in  which 

B  =  weight  of  balloon  at  t°  and  h  mm.  ; 
B1  =       "      "       "       with  vapor,  at  tt°  and  ht  mm.  ; 

v  =  capacity  of  the  balloon  in  cubic  centimeters; 
0.001293  =  weight  of  1  cc,  air  at  0°  and  760  mm.  ; 
p  =  weight  of  air  in  balloon  at  t°  and  h  mm. 


FIG.  40. 


FIG.  41. 


METHOD  OF  VICTOR  MEYER. 

Experiment  62.— In   this  method    a    known 
weight  of  substance  is  converted  into  vapor,  and 
the  volume  of  vapor  formed   is   determined  by 
measuring  the  volume  of  air  which  it  displaces. 
The  apparatus   consists   of  an  outer  cylindrical 
vessel  A,  Fig.  42,  and  an  inner  vessel  B,  which 
is  connected  with  a  tube  0.     The  vessel  B  has  a 
capacity   of    about    100    cc.,  and   is  about   200 
mm.  long.     The  tube  C,  with  its  funnel-shaped 
end  E,  is  about  600  mm.  long.     First,  a  small 
quantity  of  some  substance  with  a  boiling-point     ^, 
high  enough  to  secure  the  complete  conversion       FIG  42. 
into  vapor  of  the  substance  to  be  studied,  is   put   in   the 
bottom  of  the  vessel  A,  and  a  little  ignited  asbestos  or  dry 


OZONE.  767 

mercury  in  the  bottom  of  the  vessel  B.  The  substance  in  A 
is  now  heated  to  boiling,  and  E  is  closed  with  a  rubber  stop- 
per. After  a  time  the  temperature  of  the  air  in  B  is  raised  to 
that  of  the  vapor  in  A,  and  no  more  escapes  from  the  tube  D. 
When  this  condition  of  equilibrium  is  reached,  a  small  weighed 
quantity  of  the  substance  under  examination  is  dropped  into 
the  vessel  B,  the  stopper  being  removed  from  E  and  quickly 
replaced.  The  substance  is  converted  into  vapor,  and  displaces 
an  equivalent  volume  of  air,  and  this  displaced  air  is  collected 
over  water  in  the  measuring-tube  placed  over  the  end  of  D. 
When  no  more  air  escapes,  the  volume,  is  determined  in  the 
usual  way.  The  specific  gravity  of  the  substance  is  calculated 
by  the  aid  of  the  following  formula: 

£  _  s,     (1  4-  0.00366  X  Q760 
"      '  (B  —  w)Vx  0.001293' 

in  which  0.001293  is  the  weight  of  1  cc.  air  in  grams  at  760 
mm.  and  0°;  and,  further, 

S  =  weight  of  substance  taken; 

t  =  temperature  of  the  room,  or  of  the  water  in  the  measur- 
ing apparatus; 
B  =  height  of  barometer; 
w  =  tension  of  aqueous  vapor; 
V  =  observed  volume  of  air; 

or,  the  formula  can  be  simplified  by  division,  when  it  takes 
this  form: 

(1  +  0.00366  X  0587,780 
(B-w)V 

The  above  is  the  simplest  form  of  apparatus  used.  To  avoid 
opening  and  shutting  the  vessel  in  order  to  introduce  the  sub- 
stance, an  arrangement  has  been  devised  for  holding  the  sub- 
stance below  the  stopper,  until  the  proper  temperature  is 
reached,  and  then  releasing  it  without  disturbing  the  stopper. 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  VII. 
OZONE. 

Experiment  63. — Put  a  few  sticks  of  ordinary  phosphorus 
on  the  bottom  of  a  good-sized  bottle  with  a  wide  mouth,  and 


768    EXPERIMENTS  TO  ACCOMPANY  CHAPTER   VIII. 

partly  cover  the  phosphorus  with  water.  In  a  short  time  the 
odor  of  ozone  will  be  perceptible,  and  the  gas  can  also  be  de- 
tected by  means  of  strips  of  paper  which  have  been  moistened 
with  a  dilute  solution  of  potassium  iodide  and  starch-paste. 
See  whether  such  papers  are  changed  in  the  air  ?  What  is  the 
cause  of  the  change?  If  convenient,  examine  the  air  in  the 
neighborhood  of  a  frictional  electrical  machine,  and  see 
whether  it  causes  the  papers  to  change  color. 

HYDROGEN  DIOXIDE. 

Experiment  64. — Finely  powder  some  barium  dioxide,  and 
add  some  of  it  to  dilute  sulphuric  acid.  Filter  from  the  pre- 
cipitated barium  sulphate,  and  with  the  solution  try  the  fol- 
lowing reactions : 

Heat  some  in  a  test-tube.  What  takes  place? — Add  to  an- 
other small  portion  a  little  of  a  dilute  solution  of  potassium 
permanganate.  To  another  portion  add  a  little  finely  pow- 
dered manganese  dioxide.  What  is  given  off  ? — To  a  dilute 
solution  contained  in  a  small  stoppered  cylinder  add  a  few 
drops  of  a  dilute  solution  of  potassium  dichromate,  and  quick- 
ly add  ether,  and  shake  the  cylinder  thoroughly. 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  VIII. 
PREPARATION  OF  CHLORINE. 

Experiment  65. — Pour  2  or  3  cc.  concentrated  sulphuric 
acid  on  a  gram  or  two  of  common  salt  in  a  test-tube.  A  gas 
will  be  given  off  which  forms  dense  white  fumes  in  the  air  and 
has  a  sharp,  penetrating  taste  and  smell.  This  is  hydrochloric 
acid  gas. 

Experiment  66. — Pour  2  or  3  cc.  concentrated  sulphuric  acid 
on  a  few  grams  of  manganese  dioxide  in  a  test-tube.  Heat, 
and  examine  the  gas  given  off.  jQonvince  yourself  that  it  is 
oxygen. 

Experiment  67. — Mix  2  grams  manganese  dioxide  and  2 
grams  common  salt.  Pour  4  to  5  cc.  dilute  sulphuric  acid  on 
the  mixture  in  a  test-tube.  This  experiment  should  be  per- 
formed under  a  hood  in  which  the  draught  is  good,  as  the  gas 
which  is  given  off  is  not  only  disagreeable,  but  irritating  to 
the  respiratory  organs.  Notice  the  color  and  odor  of  the  gas. 
[Does  it  support  combustion  ?  Does  it  burn  ?] 


PREPARATION  OF  CHLORINE.  ?69 

The  best  way  to  make  chlorine  is  the  following :  Mix  5  parts 
coarsely  granulated  manganese  dioxide  and  5  parts  coai 
ly  granulated  common  salt. 
Make  a  mixture  of  12  parts 
concentrated  sulphuric  acid 
and  6  parts  water.  Let  this 
mixture  cool  down  to  the  tem- 
perature of  the  room,  and  then 
pour  it  upon  the  mixture  of 
salt  and  manganese  dioxide. 
Gently  heat  on  a  sand-bath, 
and  a  regular  current  of  chlo- 
rine will  be  given  off.  The  gas 
is  collected  by  displacement  of 
air  in  a  dry  glass  vessel.  The 
apparatus  for  the  purpose  is 
arranged  as  shown  in  Fier.  43. 

FIG  43. 

The  delivery-tube  should  reach 

to  the  bottom  of  the  collecting  vessel,  and  the  mouth  of  the 
vessel  should  be  covered  with  a  piece  of  paper  to  prevent  cur- 
rents of  air  from  carrying  away  the  chlorine.  As  the  gas  col- 
lects in  the  vessel  the  experimenter  can  judge  of  the  quantity 
present  by  means  of  the  color. 

Experiment  68. — Collect  six  or  eight  dry  cylinders  or  bot- 
tles full  of  chlorine.  Make  the  gas  from  about  30  grams  of 
manganese  dioxide,  using  the  other  substances  in  the  propor- 
tions already  stated. 

(1)  Introduce  into  one  of  the  vessels  containing  chlorine  a 
little  finely  powdered  antimony. 

(2)  Into  a  second  vessel  put  a  few  pieces  of  heated  thin 
copper- foil. 

(3)  Into  a  third  vessel  put  a  piece  of  paper  with  some  writ- 
ing on  it,  some  flowers,  and  pieces  of  cotton  print.     The  sub- 
stances used  must  be  moist. 

(4)  Into  a  fourth  vessel  put  a  dry  piece  of  the  same  cotton 
print  as  that  used  in  the  previous  experiment. 

What  conclusions  do  the  results  of  the  above  experiments 
justify  as  to  the  conduct  of  chlorine? 

Experiment  69. — Cut  a  piece  of  filter-paper  about  an  inch 
wide  and  six  to  eight  inches  long.  Pour  on  this  some  ordinary 
oil  of  turpentine  previously  warmed  slightly.  Introduce  this 
into  one  of  the  vessels  of  chlorine.  A  flash  of  flarne  is  noticed, 


and  a  dense  black  cloud  is  formed.  The  action  in  this  case 
is  due  to  the  great  affinity  of  chlorine  for  hydrogen.  Oil  of 
turpentine  consists  of  carbon  and  hydrogen.  The  main  action 
of  the  chlorine  consists  in  extracting  the  hydrogen  and  leaving 
the  carbon.  The  experiment  is  interesting  chiefly  in  so  far 
as  it  illustrates  the  general  tendency  of  chlorine  to  act  upon 
vegetable  substances. 

CHLORINE  DECOMPOSES  WATER  IN  THE  SUNLIGHT. 

Experiment  70. — Seal  the  end  of  a  glass  tube  about  a  metre 
(or  about  a  yard)  long  and  about  12  mm.  (|  inch) 
internal  diameter.  Fill  this  with  a  strong  solu- 
tion of  chlorine  in  water.  Invert  it  as  shown  in 
Fig.  44,  in  a  shallow  vessel  containing  some  of 
the  same  solution  of  chlorine  in  water.  Place 
the  tube  in  direct  sunlight.  Gradually  bubbles 
of  gas  will  be  seen  to  rise  and  collect  in  the  up- 
per end,  and  the  color  of  the  solution,  which  is 
at  first  greenish  yellow,  like  that  of  chlorine, 
disappears.  The  gas  can  be  shown  to  be  oxygen. 

CHLORINE  HYDRATE. 

Experiment  71. — Conduct  chlorine  into  a  flask 
containing  water  cooled  down  to  about  2°  or  3° 
Centigrade.     If  crystals  are  formed  remove  some 
by  filtering  out-of-doors  if  the  weather  is  cold.     Expose  some 
of  the  crystals  on  filter-paper  under  a  hood  in  the  laboratory. 
What  changes  have  taken  place  ? 

FORMATION  OF  HYDROCHLORIC  ACID. 

Experiment  72.— Light  a  jet  of  hydrogen  in  the  air  and 
carefully  introduce  it  into  a  vessel  containing  chlorine.  It 
will  continue  to  burn,  but  the  flame  will  not  appear  the  same. 
A  gas  will  be  given  off  which  forms  clouds  in  the  air.  This 
gas  has  a  sharp,  penetrating  taste  and  smell. 

Experiment  73.— Half  fill  a  small,  wide-mouthed  cylinder 
over  hot  water  with  chlorine  gas.  Then  fill  it  with  hydrogen. 
The  direct  sunlight  must  not  shine  upon  the  cylinder  while 
it  contains  the  mixture.  Turn  it  mouth  upward  and  apply 
a  flame. 


PREPARATION  OF  HYDROCHLORIC  ACID.  771 

PREPARATION  OF  HYDROCHLORIC  ACID. 

- — 

Experiment  74.— Arrange  an  apparatus  as  shown  in  Fig.  45i 


FIG.  45. 

"Weigh  out  5  parts  common  salt,  5  parts  concentrated  sul- 
phuric acid,  and  1  part  water.  Mix  the  acid  and  water,  tak- 
ing the  usual  precautions;  let  the  mixture  cool  down  to  the 
ordinary  temperature,  and  then  pour  it  on  the  salt  in  the 
flask.  For  the  purposes  of  the  experiment  take  about  20 
grams  of  salt.  Now  heat  the  flask  gently,  and  the  gas  will 
be  regularly  evolved.  Conduct  it  at  first  through  water  con- 
tained in  the  two  Wolff's  bottles  until  what  passes  over  is 
all  absorbed  in  the  first  bottle.  The  reason  why  gas  at  first 
bubbles  through  all  the  bottles  is,  that  the  apparatus  is  full 
of  air,  which  is  first  driven  out.  When  the  air  has  been  dis- 
placed, the  gas  is  all  absorbed  as  soon  as  it  comes  in  contact 
with  the  water. — After  the  gas  has  passed  for  ten  to  fifteen 
minutes,  disconnect  at  A.  Notice  the  fumes.  These  become 
denser  by  blowing  the  breath  on  them.  Why? — Apply  a 
lighted  match  to  the  end  of  the  tube.  Does  the  gas  burn? 
— Collect  some  of  the  gas  in  a  dry  cylinder  by  displacement 
of  air,  as  in  the  case  of  chlorine.  The  specific  gravity  of  the 
gas  being  1.26,  the  vessel  must  of  course  be  placed  with  the 
mouth  upward.  That  the  gas  is  colorless  and  transparent  is 
shown  by  the  appearance  of  the  generating  flask,  which  is 
filled  with  the  gas.  Insert  a  burning  stick  or  candle  in  the 
cylinder  filled  with  the  gasv — Reconnect  the  gen erat ing-flask 
with  the  series  of  bottles  containing  water,  and  let  the  pro- 
cess continue  until  no  more  gas  comes  over.  The  reaction 
represented  in  the  equation 

H2S04  =  Na2S04  +  2HC1 


772      EXPERIMENTS  TO  ACCOMPANY  CHAPTER  IX. 

is  now  complete.  Disconnect  the  flask,  and  after  it  has 
cooled  down  pour  water  on  the  contents  ;  when  the  substance 
is  dissolved  filter  it  and  evaporate  to  such  a  concentration 
that,  on  cooling,  the  sodium  sulphate  is  deposited.  Pour  off 
the  liquid  and  dry  the  solid  substance  by  means  of  filter- 
paper.  Compare  the  substance  with  the  common  salt  which, 
you  put  in  the  flask  before  the  experiment.  What  proofs 
have  you  that  the  two  substances  are  not  the  same? — Heat 
a  small  piece  of  each  in  a  dry  tube  closed  at  one  end.  What 
differences  do  you  notice  ? — Treat  a  small  piece  of  each  in  a 
test-tube  with  sulphuric  acid.  What  difference  do  you  no- 
tice ? — If  in  the  experiment  we  should  recover  all  the  sodium 
sulphate  formed,  how  much  should  we  have  ? — Put  about  50 
cc.  of  the  liquid  from  the  first  Wolff's  bottle  in  a  porcelain 
evaporating-dish.  Heat  over  a  small  flame  just  to  boiling.  Is 
hydrochloric  acid  given  off  ?  Can  all  the  liquid  be  driven  off 
by  boiling  ? — Try  the  action  of  the  solution  on  some  iron 
filings.  What  is  given  off  ? — Add  some  to  a  little  granulated 
zinc  in  a  test-tube.  What  is  given  off  ? — Add  a  little  to 
some  manganese  dioxide  in  a  test-tube.  What  is  given  off  ? 
— Add  ten  or  twelve  drops  of  the  acid  to  2  to  3  cc.  water  in  a 
test-tube.  Taste  the  dilute  solution.  It  has  what  is  called  a 
sour  or  acid  taste,  the  two  terms  being  practically  synony- 
mous.— Add  a  drop  or  two  of  a  solution  of  blue  litmus, 
or  put  into  it  a  piece  of  paper  colored  blue  with  litmus. 
What  change  takes  place?  Litmus  is  a  vegetable  color  pre- 
pared for  use  as  a  dye.  Other  vegetable  colors  are  changed 
by  hydrochloric  acid. — Steep  a  few  leaves  of  red  cabbage  in 
water.  Add  a  few  drops  of  the  solution  thus  obtained  to  di- 
lute hydrochloric  acid.  Is  there  any  change  in  color  ? — The 
color  will  be  restored  in  each  case  by  adding  a  few  drops  of  a 
solution  of  caustic  soda. — In  what  experiment  has  caustic 
soda  been  obtained  ?  What  relation  does  it  bear  to  water  ? — 
To  the  dilute  solution  of  hydrochloric  acid  add  drop  by  drop 
a  dilute  solution  of  caustic  soda.  Is  the  acid  taste  destroyed  ? 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  IX. 

CHLORIC  ACID  AKD  POTASSIUM  CHLORATE. 

Experiment  75. — Dissolve  40  grams  (or  about  1J  ounces) 
caustic  potash  in  100  cc.  water  in  a  beaker-glass,  and  pass 
chlorine  into  it.  When  chlorine  passes  freely  through  the 


PERCHLORIC  ACID. 


773 


solution,  thus  indicating  that  it  is  no  longer  absorbed,  stop 
the  action.  After  boiling  filter  the  solution  and  allow  it  to 
cool,  when  crystals  of  potassium  chlorate  will  be  deposited, 
mixed  with  a  little  potassium  chloride.  Eecrystallize  from  a 
little  water.  Filter  off  the  crystals  and  dry  them.  What  evi- 
dence have  you  that  the  substance  is  potassium  chlorate?  Does 
it  give  off  oxygen  when  heated  ?  In  a  dry  test-tube  pour  two 
or  three  drops  of  concentrated  sulphuric  acid  on  a  small  crys- 
tal of  the  substance.  Do  the  same  with  a  piece  of  potassium 
chlorate  from  the  laboratory  bottle.  Hold  the  mouth  of  the 
test-tube  away  from  the  face.  What  is  noticed  in  each  case  ? 
— Evaporate  the  solution  from  which  the  crystals  of  potassi- 
um chlorate  have  been  removed.  On  allowing  it  to  cool 
crystals  will  again  be  deposited.  Take  them  out  and  recrys- 
tallize  them.  Does  this  substance  give  off  oxygen  when 
heated  ?  Does  it  give  off  a  gas  when  treated  with  sulphuric 
acid  ?  Is  this  gas  colored  ?  Is  it  hydrochloric  acid  ?  How 
do  you  know  that  it  is  ?  If  the  gas  is  hydrochloric  acid,  what 
is  the  solid  substance  from  which  it  is  formed  ?  And  what 
is  left  in  the  test-tube  ? 

Experiment  76. — Mix  10 
grams  fresh  quick-lime  with 
20  cc.  water.  After  the  slak- 
ing is  over,  pass  chlorine  into 
it  until  the  gas  is  no  longer 
absorbed.  Put  the  powder 
thus  formed  in  a  flask  ar- 
ranged as  shown  in  Fig.  46. 
Pour  a  mixture  of  equal  parts 
of  sulphuric  acid  and  water 
slowly  through  the  funnel- 
tube.  Collect  by  displacement 
of  air  the  gas  given  off.  What 
evidence  have  you  that  the. 
gas  is  chlorine  ?  yIG.  40. 

PERCHLORIC  ACID. 

Experiment  77. — Make  potassium  perchlorate  as  follows : 
Gently  heat  50  to  100  grams  potassium  chlorate  until  after 
having  been  liquid  it  becomes  thick  and  pasty,  and  gas  is  not 
given  off  without  raising  the  temperature.  After  cooling, 
break  up  the  mass  and  treat  it  with  cold  water.  This  dissolves 


774      EXPERIMENTS  TO  ACCOMPANY  CHAPTER  X. 

out  the  potassium  chloride  and  leaves  the  perchlorate,  which 
can  then  be  crystallized  from  hot  water.  After  the  crystallized 
salt  is  dried  it  is  decomposed  by  sulphuric  acid.  To  effect 
this  decomposition,  the  finely  powdered  salt  (10  parts)  is 
treated  in  a  retort  with  20  parts  of  pure  sulphuric  acid  which 
is  free  from  nitric  acid  and  diluted  with  -fa  its  volume  of  water. 
The  retort  is  connected  with  a  receiver  which  can  be  well 
cooled.  The  mixture  is  heated,  and  when  the  perchloric  acid 
begins  to  come  over,  the  heat  is  so  regulated  that  the  tem- 
perature does  not  rise  above  140°.  When  the  mixture  has 
become  colorless  the  operation  is  ended. 


EXPERIMENTS  TO  ACCOMPANY  CHAPTER  X. 

NEUTRALIZATION  OF  ACIDS  AND  BASES  ;  FORMATION  OF 

SALTS. 

Experiment  78. — Make  dilute  solutions  of  nitric,  hydro- 
chloric,  and  sulphuric  acids  (1  part  dilute  acid,  such  as  is  used 
in  the  laboratory,  to  50  parts  water), 
and  of  caustic  soda  and  caustic  pot- 
ash (about  1  ^ram  to  200  cc.  of 
water).  Measure  off  about  20  cc.  of 
one  of  the  acid  solutions.  Add  a 
few  drops  of  a  solution  of  blue 
litmus.  Gradually  add  to  the  meas- 
ured quantity  of  acid  sufficient  di- 
lute caustic  soda  to  cause  the  red 
color  just  to  change  to  blue.  As  long 
as  the  solution  is  red  it  is  acid. 
"When  it  turns  blue  it  is  alkaline. 
At  the  turning-point  it  is  neutral. 
The  operation  is  best  carried  on  by 
means  of  a  burette,  which  is  a  gradu- 
ated tube  with  an  opening  from 
which  small  quantities  can  be  poured. 
A  convenient  shape  is  that  repre- 
sented in  Fig.  47.  At  the  lower  end 
is  a  small  opening.  The  flow  of  the 
•liquid  from  the  burette  is  controlled 
by  means  of  a  small  pinch-cock.  It 
will  require  some  practice  to  enable  the  student  to  know  ex- 


FIG.  47. 


STUDY  OF  THE  PRODUCTS  FORMED.  775 

actly  when  the  red  color  disappears  and  the  blue  appears,  but 
with  practice  the  point  can  be  discerned  with  great  accuracy. 
Should  too  much  alkali  be  allowed  to  get  into  the  acid,  add  a 
small  measured  quantity  of  the  acid  from  another  burette. 
Having  in  one  experiment  determined  how  much  of  the  solu- 
tion of  alkali  is  required  to  cause  the  red  color  to  change  to 
blue  in  operating  on  a  given  quantity  of  the  acid  solution,  try 
the  experiment  again,  using  a  different  quantity  of  the  acid 
solution.  If  the  results  of  several  experiments  with  the  same 
acid  and  alkali  are  recorded,  it  will  be  found  that  there  is 
a  definite  ratio  between  the  quantities  of  acid  and  alkali  so- 
lution required  to  neutralize  one  another.  If,  for  example, 
15  cc.  of  the  alkali  solution  are  required  to  neutralize  20  cc.  of 
the  acid  solution,  18  cc.  of  the  alkali  solution  will  be  required 
to  neutralize  24  cc.  of  the  acid  solution,  30  cc.  to  neutralize 
40  cc.,  etc.  In  other  words,  in  order  to  neutralize  a  given 
quantity  of  an  acid,  a  definite  quantity  of  an  alkali  is  necessary. 
Perform  similar  experiments  with  the  other  acids.  Afterwards 
carefully  examine  the  numerical  results.  Suppose  it  should 
require  15  cc.  of  the  caustic-soda  solution  or  12  cc.  of  the 
caustic-potash  solution  to  neutralize  20  cc.  of  the  hydrochloric- 
acid  solution.  Compare  the  quantities  of  these  alkali  solu- 
tions necessary  to  neutralize  equal  quantities  of  the  other  acids. 
What  conclusion  is  justified  with  reference  to  the  act  of  neu- 
tralization ? 

STUDY  OF  THE  PKODUCTS  FOBMED. 

Experiment  79. — Dissolve  about  10  grams  caustic  soda  in 
100  cc.  water.  Add  hydrochloric  acid  slowly,  examining  the 
solution  from  time  to  time  by  means  of  a  piece  of  paper  col- 
ored blue  with  litmus.  As  long  as  the  solution  is  alkaline  it 
will  cause  no  change  in  the  color  of  the  paper.  The  instant 
the  point  of  neutralization  is  passed,  the  solution  changes  the 
color  of  the  paper  to  red;  when  exactly  neutral,  it  will  neither 
change  the  blue  to  red,  nor,  if  the  color  is  changed  to  red 
by  means  of  another  acid,  will  it  change  it  back  again.  When 
this  point  is  reached,  evaporate  to  complete  dryness  on  the 
water-bath,  and  see  what  is  left.  Taste  the  substance.  Has 
it  an  acid  taste?  Does  it  suggest  any  familiar  substance?  If 
it  is  sodium  chloride,  how  ought  it  to  conduct  itself  when 
treated  with  sulphuric  acid?  Does  it  conduct  itself  in  this 
way  ?  Satisfactory  evidence  can  be  given  that  the  substance 


776     EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XII. 

is  sodium  chloride.  It  is  not  an  acid  nor  an  alkali.  It  is 
neutral. 

Experiment  80. — Perform  a  similar  experiment,  using 
dilute  nitric  acid  and  caustic  soda.  What  evidence  have  you 
that  the  product  in  this  case  is  different  from  caustic  soda  ? 

Experiment  81. — Perform  similar  experiments  with  dilute 
sulphuric  acid  and  caustic  soda;  with  sulphuric  acid  and 
caustic  potash;  with  nitric  acid  and  caustic  potash;  with  hy- 
drochloric acid  and  caustic  potash.  Dry  and  examine  the 
product  carefully  in  each  case;  and  keep  for  future  study  what 
is  not  used  in  these  experiments. 

FOR  CHAPTER  XI. 

A  large  table  of  the  Natural  System  of  the  Elements,  like 
that  on  page  151,  should  be  hung  up  in  a  conspicuous  place 
in  the  laboratory.  It  would  be  well  also  to  have  such  a  table 
pasted  upon  a  cylinder  which  can  be  revolved  on  its  axis,  so 
that  the  continuity  of  the  system  may  be  impressed  upon  the 
mind. 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XII. 
PREPARATION  OF  BROMINE. 

Experiment  82. — Mix  together  3.5  grams  potassium  bro- 
mide and  7  grams  manganese  dioxide.  Put  the  mixture  into  a 
500  cc.  flask;  connect  with  a  condenser  (see  Fig.  39).  Mix  15 
cc.  concentrated  sulphuric  acid  and  90  cc.  water.  After  cool- 
ing pour  the  liquid  on  the  mixture  in  the  flask.  Gently  heat, 
when  bromine  will  be  given  off  in  the  form  of  vapor.  A  part 
of  this  will  condense  and  collect  in  the  receiver.  Perform  this 
experiment  under  a  hood  with  a  good  draught. 

HYDROBROMIC  ACID. 

Experiment  83. — In  a  small  porcelain  evaporating-dish 
put  a  few  crystals,  of  potassium  bromide.  Pour  on  them  a 
few  drops  of  concentrated  sulphuric  acid.  The  white  fumes 
of  hydrobromic  acid  and  the  reddish-brown  vapor  of  bromine 
are  noticed.  Treat  a  few  crystals  of  potassium  or  sodium 
chloride  in  the  same  way.  What  difference  is  there  between 
the  two  cases  ? 

The  preparation  of  hydrobromic  acid  may  be  shown  in  the 
lecture-room  as  follows: 


HTDROBROMIO  ACID.  777 

Experiment  84. — Arrange  an  apparatus  as  shown  in  Fig. 
48.     In  the  flask  put  1  part  red  phosphorus  and  2  parts  water. 


FIG.  48. 

Let  10  parts  bromine  gradually  drop  into  the  flask  from  the 
glass-stoppered  funnel.  Pass  the  gas  through  a  U-tube 
loosely  packed  with  asbestos  containing  red  phosphorus  in 
order  to  free  the  hydrobromic  acid  from  bromine,  which  to 
some  extent  passes  over  with  it.  Collect  some  of  the  gas  in 
water,  and  examine  the  solution.  How  does  the  gas  act  when 
allowed  to  escape  in  the  air  ?  Fill  a  cylinder  with  the  gas  in 
the  same  way  as  was  done  with  hydrochloric  acid,  and  fill  an- 
other with  chlorine.  While  covered  with  glass  plates  bring 
their  mouths  together.  Then  withdraw  the  plates.  What 
change  is  observed  ?  What  is  this  due  to  ? 

Experiment  85. — To  a  dilute  solution  of  sodium  hydroxide 
add  bromine  water  made  by  shaking  up  a  little  liquid  bromine 
in  a  bottle  with  water.  What  change  takes  place  ?  Add  sul- 
phuric acid  until  the  liquid  shows  an  acid  reaction.  What 
takes  place?  The  changes  here  referred  to  are  perfectly  anal- 
ogous to  those  which  would  *take  place  if  chlorine  were  used 
instead  of  bromine.  Shake  a  solution  containing  a  little  free 
bromine  with  ether  ;  with  chloroform  ;  with  carbon  disulphide. 
What  changes  do  you  observe  ? 


778    EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XII. 


IODINE. 

Experiment  86. — Mix  about  2  grams  of  sodium  or  potas- 
sium iodide  and  4  grams  manganese  dioxide.  Treat  with  a 
little  concentrated  sulphuric  acid  in  a  one  to  two  liter  flask. 
Heat  gently  on  a  sand-bath.  Gradually  the  vessel  will  be  filled 
with  the  beautiful  colored  vapor  of  iodine.  In  the  upper  parts 
of  the  flask  some  of  the  iodine  will  be  deposited  in  the  form 
of  crystals  of  a  grayish-black  color. 

Experiment  87. — Make  solutions  of  iodine  in  water,  in 
alcohol,  and  in  a  water  solution  of  potassium  iodide.  Use 
small  quantities  in  test-tubes. 

Experiment  88. — Dissolve  a  piece  of  potassium  iodide  the 
size  of  a  small  pea  in  about  100  cc.  water  in  a  stoppered  cylin- 
der. Add  enough  carbon  disulphide  to  make  a  layer  about 
an  inch  thick  at  the  bottom  of  the  cylinder.  Shake  the  two 
liquids  together.  Does  the  carbon  disulphide  become  colored  ? 
Add  a  drop  of  chlorine  water  and  shake  again.  What  differ- 
ence do  you  observe  in  the  two  cases  ?  Explain  this.  Try 
the  same  experiment,  using  chloroform  instead  of  carbon 
disulphide. 

IODINE  CAN  BE  DETECTED  BY  MEANS  OF  ITS  ACTION  UPON 
STAECH-PASTE. 

Experiment  89. — Make  some  starch-paste  by  covering  a 
few  grains  of  starch  in  a  porcelain  evaporating-dish  with  cold 
water,  grinding  this  to  a  paste,  and  pouring  200-300  cc.  boil- 
ing-hot water  on  it.  After  cooling  add  a  little  of  this  paste 
to  a  dilute  water  solution  of  iodine.  The  solution  will  turn 
blue  if  the  conditions  are  right.  Now  add  a  little  of  the  paste 
to  a  diluted  water  solution  of  potassium  iodide.  Is  there  any 
change?  Add  a  drop  or  two  of  a  solution  of  chlorine  in  water. 
Why  the  difference  ?  Will  not  chlorine  water  alone  act  this 
way  toward  starch-paste  ? 

ACTION  OF  SULPHURIC  ACID  UPON  POTASSIUM  IODIDE. 

Experiment  90. — Bring  a  piece  of  potassium  iodide  the 
size  of  a  pea  in  a  dry  test-tube  ;  add  one  drop  of  water  and 
three  or  four  drops  of  concentrated  sulphuric  acid  ;  the  salt 
becomes  brown  ;  heat  gently  ;  violet-colored  vapor  escapes, 
and  with  it  a  gas  with  an  odor  like  that  of  rotten  eggs.  At 


IODIC  ACID— PROPERTIES  OF  SULPHUR.  779 

the  same  time  a  yellow  coating  appears  on  the  inside  of  the 
tube  above  the  acid.  Add  five  or  six  drops  more  of  the  acid 
and  continue  to  heat  gently.  The  bad  odor  first  noticed  dis- 
appears gradually,  and  another,  quite  different  odor,  irritating 
to  the  throat  is  now  perceptible.  This  is  sulphur  dioxide, 

so, 

IODIC  ACID. 

Experiment  91. — Pass  chlorine  into  a  test-tube  containing 
iodine  in  suspension  in  water  ;  or  add  chlorine  water.  What 
becomes  of  the  iodine? 

Experiment  92. — Add  chlorine  water  to  a  dilute  solution 
of  potassium  iodide,  and  note  the  successive  changes. 

Experiment  93. — Dissolve  iodine  in  caustic  soda.  Add 
an  acid  to  the  solution.  Explain  the  changes. 

HYDROFLUORIC  ACID. 

Experiment  94. — In  a  lead  or  platinum  vessel  put  a  few 
grams  (5-6)  of  powdered  fluor-spar  and  pour  on  it  enough 
concentrated  sulphuric  acid  to  make  a  thick  paste.  Cover  the 
surface  of  a  piece  of  glass  with  a  thin  layer  of  wax  or  paraffin, 
and  through  this  scratch  some  letters  or  figures,  so  as  to  leave 
the  glass  exposed  where  the  scratches  are  made.  Put  the 
glass  over  the  vessel  containing  the  fluor-spar,  and  let  it  stand 
for  some  hours.  Take  off  the  glass,  scrape  off  the  coating, 
and  the  figures  which  were  marked  through  the  wax  or  paraf- 
fin will  be  found  etched  on  the  glass. 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XIII. 
PROPERTIES  OF  SULPHUR. 

Experiment  95. — Distil  about  10  grams  roll  sulphur  from 
an  ordinary  glass  retort.  What  changes  in  color  and  in  con- 
dition take  place  ?  Collect  the  liquid  sulphur  formed  by  the 
condensation  of  the  vapor  in  a  beaker-glass  containing  cold 
water. 

Experiment  96. — Treat  some  powdered  roll  sulphur  with 
carbon  disulphide  and  filter.  Does  it  all  dissolve  ?  Try  the 
same  experiment  with  flowers  of  sulphur.  Does  this  all  dis- 
solve? Put  the  solutions  together  and  allow  to  evaporate. 
Examine  the  crystals  deposited.  Compare  them  with  some 
natural  crystals  of  sulphur.  See  whether  one  of  the  crystals 
will  completely  dissolve  in  carbon  disulphide. 


rtfO    EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XIII. 

Experiment  97. — In  a  covered  sand  or  Hessian  crucible 
melt  about  25  grams  of  roll  sulphur.  Let  it  cool  slowly,  and 
when  a  thin  crust  has  formed  on  the  surface  make  a  hole 
through  this  and  pour  out  the  liquid  part  of  the  sulphur. 
What  is  left  ?  Compare  with  the  crystals  formed  in  the  last 
experiment. — Lay  the  crucible  aside,  and  in  the  course  of  a 
few  days  again  examine  the  crystals.  What  changes,  if  any, 
have  taken  place? 

Experiment  98. — Add  hydrochloric  acid  to  a  solution  of 
sodium  thiosulphate.  What  takes  place  ? 

Experiment  99. — In  a  wide  test-tube  heat  some  sulphur 
to  boiling.  Introduce  into  it  small  pieces  of  copper-foil  or 
sheet  copper.  Or  hold  a  narrow  piece  of  sheet  copper  so  that 
the  end  just  dips  into  the  boiling  sulphur. 

Experiment  100. — Dissolve  some  sulphur  in  concentrated 
caustic  soda.  In  what  form  is  the  sulphur  in  the  solution  ? 

HYDROGEN  SULPHIDE. 

Experiment  101. — Arrange  an  apparatus  as  shown  in  Fig. 
49.  Put  a  small  handful  of  the  sulphide  of  iron,  FeS,  in  the 


FIG.  49. 

flask,  and  pour  dilute  sulphuric  acid  upon  it.  Pass  the 
evolved  gas  through  a  little  water  contained  in  the  wash  cylin- 
der A.  Pass  some  of  the  gas  into  water.  [What  evidence 
have  you  that  it  dissolves  ?]  Collect  some  by  displacement  of 
air.  Its  specific  gravity  is  1.178.  Set  fire  to  some  of  the  gas 
contained  in  a  cylinder.  In  this  case  the  air  has  not  free 


MANUFACTURE  OF  SULPHURIC  ACID.  781 

access  to  the  gas,  and  the  combustion  is  not  complete.  The 
hydrogen  burns  to  form  water,  while  a  part  of  the  sulphur  is. 
deposited  upon  the  inside  walls  of  the  cylinder.  If  there  is 
free  access  of  air,  the  sulphur  burns  to  sulphur  dioxide  and 
the  hydrogen  to  water. 

Make  a  solution  of  the  gas  in  water  in  the  usual  way.  Put 
some  of  this  in  a  bottle  and  set  it  aside,  and  in  the  course  of 
a  few  days  examine  it  again.  Boil  another  portion  for  a  time 
in  a  test-tube,  and  note  the  changes.  Pass  a  little  of  the  gas 
through  concentrated  sulphuric  acid  contained  in  a  test-tube, 
and  note  the  changes.  Moisten  strips  of  paper  with  dilute 
solutions  of  lead  nitrate,  copper  sulphate,  stannous  chloride, 
.antimony  chloride,  and  mercuric  chloride  ;  and  expose  these 
papers  in  turn  to  the  gas.  What  changes  take  place  ?  Eepeat 
Experiment  90,  and  see  whether  one  of  the  gases  given  off 
produces  similar  changes. 

Experiment  102. — Pass  hydrogen  sulphide  successively 
through  solutions  containing  a  little  lead  nitrate,  cadmium 
nitrate,  and  arsenic  prepared  by  dissolving  a  little  white 
arsenic,  or  arsenic  trioxide,  As203,  in  dilute  hydrochloric  acid. 
What,  action  takes  place  in  each  case  ?  The  formula  of  lead 
nitrate  is  Pb(N03)2 ;  that  of  cadmium  nitrate,  Cd(N03)2 ;  and 
that  of  the  chloride  of  arsenic  in  solution  is  As013.  The  cor- 
responding sulphides  are  represented  by  the  formulas  PbS, 
€dS,  and  As2S3. 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XIV. 
MANUFACTURE  OF  SULPHURIC  ACID. 

Experiment  103. — The  manufacture  of  sulphuric  acid  can 
be  illustrated  in  the  laboratory  by  means  of  the  apparatus 
represented  in  Fig.  50.  This  consists  of  a  large  balloon  flask 
fitted  with  a  stopper  having  five  openings.  By  means  of  tubes 
it  is  connected  with  three  small  flasks.  One  of  these,  a,  con- 
tains water  for  the  purpose  of  providing  a  current  of  steam  ; 
another,  c,  contains  copper-foil  and  concentrated  sulphuric 
acid,  which  give  sulphur  dioxide  when  heated  ;  and  the  third, 
b,  contains  copper-foil  and  dilute  nitric  acid,  which  give  oxides 
of  nitrogen,  mainly  nitric  oxide,  NO.  When  the  nitric  oxide 
comes  in  contact  with  the  air  it  combines  with  oxygen,  form- 
ing nitrogen  trioxide  and  nitrogen  peroxide;  and  when  steam 
and  sulphur  dioxide  are  admitted  to  the  flask  the  reactions 


782    EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XIV. 

involved  in  the  manufacture  of  sulphuric  acid  take  place. 
By  means  of  a  pair  of  bellows  attached  at  d  air  is  supplied. 
If  air  is  not  forced  in,  the  gases  become  colorless,  owing  to 


FIG.  50. 

complete  reduction  of  the  oxides  of  nitrogen  to  the  form  of 
nitric  oxide,  NO,  which  is  colorless.  If  steam  is  not  admitted 
the  walls  of  the  vessel  become  covered  with  crystals  of  nitro- 
syl-sulphuric  acid.  This  is,  however,  decomposed  by  an  excess 
of  steam. 

Experiment  104. — Into  a  vessel  containing  ordinary  con- 
centrated sulphuric  acid  introduce  small  sticks  of  wood,  pieces 
of  paper,  and  various  other  organic  substances,  and  note  the 
result.  The  charring  effect  is  particularly  well  shown  by 
adding  the  acid  drop  by  drop  to  a  concentrated  solution  of 
sugar,  or  to  molasses,  and  stirring. 

Experiment  105. — Sulphuric  acid  is  detected  in  analysis 
by  adding  barium  cloride  to  its  solution,  when  insoluble  bar- 
ium sulphate  is  formed. 

H2S04  +  Bad,  =  BaS04  +  2HC1. 

Other  insoluble  sulphates  are  those  of  strontium  and  lead  ; 
and  calcium  sulphate  is  difficultly  soluble.  To  a  dilute  solu- 
tion of  sulphuric  acid  or  of  any  soluble  sulphate,  add  in  test- 
tubes  barium  chloride,  strontium  nitrate,  and  lead  nitrate. 


SULPHUROUS  ACID  AND  SULPHUR  DIOXIDE.      783 


SULPHUROUS  ACID  AND  SULPHUR  DIOXIDE. 

Experiment  106. — Put  eight  or  ten  pieces  of  sheet  copper, 
one  to  two  inches  long  and  about  half  an  inch  wide,  in  a  500 
cc.  flask  ;  pour  15  to  20  cc.  concentrated  sulphuric  acid  on  it. 
On  heating,  sulphur  dioxide  will  be  evolved,  The  moment 
the  gas  begins  to  come  off,  lower  the  flame,  and  keep  it  at 
such  a  height  that  the  evolution  is  regular  and  not  too  active. 
Pass  some  of  the  gas  into  a  bottle  containing  water.  The 
solution  in  water  is  called  sulphurous  acid. 

Experiment  107. — Pass  sulphur  dioxide  into  a  moderately 
dilute  solution  of  potassium  hydroxide,  until  the  solution  is 
saturated.  What  is  then  contained  in  the  solution?  To  a 
little  of  it  add  hydrochloric  acid.  What  takes  place? 

Experiment  108. — Try  the  effect  of  heating  concentrated 
sulphuric  acid  with  charcoal,  and  with  sulphur. 

Experiment  109. — Collect  by  displacement  of  air  some  of 
the  gas  made  in  Experiment  106.  Does  it  burn  ?  or  does  it 
support  combustion? 

Experiment  110. — Pass  some  of  the  gas  through  a  bent- 
glass  tube  surrounded  by  a  freezing  mixture  of  salt  and  ice. 
Tubes  provided  with  glass  stop-cocks  are  made  for  such  pur- 


FIG.    51. 

poses.  They  generally  have  the  form  represented  in  Fig.  51. 
If  the  tube  is  taken  out  of  the  freezing  mixture,  the  liquid 
sulphur  dioxide  changes  rapidly  to  gas,  if  the  tube  is  open. 

Experiment  111. — Burn  a  little  sulphur  in  a  porcelain 
crucible  under  a  bell-jar.  Place  over  the  crucible  on  a  tripod 
some  flowers.  In  the  atmosphere  of  sulphur  dioxide  the 
flowers  will  be  bleached. 

SULPHUROUS  ACID  is  A  SEDUCING  AGENT. 

Experiment  112. — To  a  dilute  solution  of  potassium  iodide 
in  a  test-tube  gradually  add  chlorine  water  until  the  solution 


784     EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XV. 

becomes  clear  and  colorless.  Now  add  a  solution  of  sulphur- 
ous acid.  At  first  iodine  is  deposited,  but  on  further  addi- 
tion of  sulphurous  acid  it  dissolves  again.  Explain  all  the 
changes. 

SULPHUR  TRIOXIDE. 

Experiment  113. — Heat  a  little  fuming  sulphuric  acid 
gently  in  a  test-tube.  What  takes  place  ?  Put  a  little  of  the 
acid  (5-10  cc.)  in  a  small  dry  retort  provided  with  a  glass 
stopper  and  connect  with  a  dry  glass  receiver.  Heat  the  re- 
tort gently,  and  keep  the  receiver  cool.  By  means  of  a  dry 
glass  rod  take  out  some  of  the  substance  which  collects  in  the 
receiver  and  put  it  in  water.  Lay  a  little  of  it  on  a  piece  of 
wood  and  on  a  piece  of  paper. 

Experiment  114. — Prepare  finely  divided  platinum  by 
moistening  some  fine  asbestos  with  a  solution  of  platinic 
chloride  and  heating  to  redness  in  a  porcelain  crucible.  The 
substance  thus  obtained  is  known  as  platinized  asbestos, 
Now  arrange  an  apparatus  so  that  both  oxygen  and  sulphur 
dioxide  can  be  passed  together  through  a  tube  of  hard  glass 
as  represented  in  Fig.  52.  First  pass  the  two  dried  gases 


0- 


S02 


FIG.  52. 

together  through  the  empty  tube  and  heat  a  part  of  the  tube 
by  means  of  a  burner.  Is  there  any  evidence  of  combination  ? 
Now  stop  the  currents  of  the  gases,  let  the  tube  cool  down, 
and  introduce  a  small  layer  of  the  platinized  asbestos.  Pass 
the  dried  gases  over  the  heated  asbestos.  What  takes  place? 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XV. 
PREPARATION  OF  NITROGEN. 

Experiment  115. — Place  a  good-sized  stoppered  bell- jar 
over  water  in  a  pneumatic  trough.  In  the  middle  of  a  flat 
cork  about  three  inches  in  diameter  fasten  a  small  porcelain 
crucible,  and  place  this  on  the  water  in  the  trough.  Put  in 


ANALYSIS  OF  AIR. 


785 


it  a  piece  of  phosphorus  about  twice  the  size  of  a  pea,  and  set 
fire  to  it.  Quickly  place  the  bell-jar  over  it.  At  first  some 
air  will  be  driven  out  of  the  jar.  The  burning  will  continue 
for  a  short  time,  and  then  gradually  grow  less  and  less  active, 
finally  stopping.  On  cooling,  it  will  be  found  that  the 
volume  of  gas  is  less  than  four  fifths  the  original  volume,  for 
the  reason  that  some  of  the  air  was  driven  out  of  the  vessel  at 
the  beginning  of  the  experiment.  Before  removing  the 
stopper  of  the  bell-jar  see  that  the  level  of  the  liquid  outside 
is  the  same  as  that  inside.  Try  the  effect  of  introducing  suc- 
cessively several  burning  bodies  into 'the  nitrogen, — as,  for 
example,  a  candle,  a  piece  of  sulphur,  phosphorus,  etc. 

Experiment  116. — Place  a  live  mouse  in  a  trap  in  a  bell- 
jar  over  water.  When  the  oxygen  is  used  up  the  mouse  will 
die.  After  the  animal  gives  plain  signs  of  discomfort,  it  may 
be  revived  by  taking  away  the  bell-jar  and  giving  it  a  free 
supply  of  fresh  air. 

Experiment  117. — Pass  air  slowly  over  copper  contained  in' 
a  tube  heated  to  redness  and  collect 
the  gas  which  passes  through.     Does 
it  act  like  nitrogen  ? 

Experiment  118. — In  a  good-sized 
Wolff's  bottle  provided  with  a  safety- 
funnel  and  delivery-tube  as  shown  in 
Fig.  53  put  some  copper-turnings 
and  pour  upon  them  concentrated 
ammonia,  but  not  enough  to  cover 
them.  Close  the  delivery-tube  by 
means  of  a  pinch-cock;  and  let  the 
vessel  stand.  What  evidence  of  ac- 
tion is  there?  After  a  time,  force 
some  of  the  gas  out  of  the  bottle  by 
pouring  water  through  the  funnel, 
and  opening  the  delivery-tube.  Does  FIG.  53. 

the  gas  act  like  nitrogen  ? 

ANALYSIS  OF  AIR. 

Experiment  119. — Arrange  an  apparatus  as  in  Fig.  25. 
Instead  of  a  plain  tube,  use  one  graduated  into  cubic  centi- 
meters. Enclose  60  to  80  cc.  air  in  the  tube  over  water. 
Arrange  the  tube  so  that  the  level  of  the  water  inside  and 
outside  is  the  same.  Note  the  temperature  of  the  air  and  the 


786     EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XV. 

height  of  the  barometer.  Reduce  the  observed  volume  to 
standard  conditions.  Now  introduce  a  piece  of  phosphorus, 
as  in  Experiment  26,  and  allow  it  to  stand  for  twenty-four 
hours.  Draw  out  the  phosphorus.  Again  arrange  the  tube 
so  that  the  level  of  the  water  inside  is  the  same  as  that  out- 
side. Make  the  necessary  corrections  for  temperature,  pres- 
sure, and  the  tension  of  aqueous  vapor.  It  will  be  found  that 
the  volume  has  diminished  considerably,  but  that  about  four 
fifths  of  the  gas  originally  put  in  the  tube  is  still  there.  If 
the  work  is  done  properly,  the  volume  of  the  gas  left  in  the 
tube  will  be  to  the  total  volume  used  as  79  to  100.  In  other 
words,  of  every  100  cc.  air  used  21  cc.  are  absorbed  by  phos- 
phorus, and  79  cc.  are  not.  The  gas  absorbed  is  oxygen, 
identical  with  the  oxygen  made  from  the  oxide  of  mercury, 
manganese  dioxide,  and  potassium  chlorate.  The  gas  left 
over  has  no  chemical  properties  in  common  with  oxygen. 
Carefully  take  the  tube  out  of  the  vessel  of  water,  closing  its 
mouth  with  the  thumb  or  some  suitable  object  to  prevent  the 
contents  from  escaping.  Turn  it  with  the  mouth  upward,  and 
introduce  into  it  a  burning  stick.  Does  it  support  combus- 
tion ?  Is  it  oxygen  ? 

Experiment  120. — Expose  a  few  pieces  of  calcium  chlo- 
ride on  a  watch-glass  to  the  air.  It  gradually  becomes  liquid 
by  absorbing  water  from  the  air. 

Experiment  121. — Expose  some  clear  lime-water  to  the 
air.  It  soon  becomes  covered  with  a  white  crust.  A  similar 
change  takes  place  if  baryta- water  is  exposed  in  the  same  way. 
Lime-water  is  made  by  putting  a  few  pieces  of  quick-lime  in 
a  bottle  and  pouring  water  upon  it.  The  mixture  is  well 
shaken  up  and  allowed  to  stand.  The  undissolved  substance 
settles  to  the  bottom,  and  with  care  a  clear  liquid  can  be 
poured  off  the  top.  This  is  lime-water,  which  is  a  solution 
of  calcium  hydroxide,  Ca(OH)3,  in  water.  Baryta- water  is 
a  solution  of  a  similar  compound  of  the  element  barium. 
When  these  solutions  are  exposed  to  nitrogen  or  oxygen,  or  to 
an  artificially  prepared  mixture  of  the  two  gases,  no  change 
takes  place.  Further,  if  air  is  first  passed  through  a  solution 
of  caustic  soda  it  no  longer  has  the  power  to  cause  the  forma- 
tion of  a  crust  on  lime-water  or  baryta-water. 

Experiment  122. — Arrange  an  apparatus  as  shown  in  Fig. 
54.  The  wash-cylinders  A  and  B  are  half  filled  with  ordi- 
nary caustic-soda  solution.  The  bottle  C  is  filled  with  water. 


ANALYSIS  OF  AIR. 


787 


The  tube  D,  which  should  be  filled  with  water  and  provided 
with  a  pinch-cock,  acts  as  a  siphon.  Open  the  pinch-cock 
and  let  the  water  flow  slowly  out  of  the  bottle.  As  it  flows 
out  air  will  be  drawn  in  through  the  caustic  soda  in  the  wash- 


Fio.  54. 

cylinders.  When  the  bottle  is  a  quarter  filled  with  air  pour 
some  water  in  again  until  it  is  full.  Then  draw  all  the 
water  off.  Now  remove  the  stopper  from  the  bottle,  pour 
in  20  to  30  cc.  lime-water  and  cork  the  bottle.  The  crust 
formed  on  the  lime-water  will  now  be  hardly,  if  at  all,  per- 
ceptible. There  is,  therefore,  something  present  in  the  air 
under  ordinary  circumstances  which  has  the  power  to  form  a 
crust  on  lime-water  or  baryta-water,  and  which  can  be  re- 
moved by  passing  the  air  through  caustic  soda.  Thorough 
examination  has  shown  that  this  is  the  compound  which 
chemists  call  carlon  dioxide,  and  which  is  commonly  known  as 
carbonic  acid  gas.  It  is  the  substance  which  was  obtained  by 
burning  charcoal  in  oxygen. 

Experiment  123. — Into  the  bottle  containing  the  air  from 
which  the  carbon  dioxide  has  been  removed  hold  a  burning 
stick  or  taper  for  a  moment.  Notice  whether  a  crust  is  now 
formed  on  the  lime-water.  Wood  and  the  material  from 
which  the  taper  is  made  contain  carbon.  Explain  the  forma- 
tion of  the  crust  on  the  lime-water  after  the  stick  of  wood  or 
taper  has  burned  for  a  short  time  in  the  vessel. 

Experiment  124. — Arrange  an  apparatus  as  shown  in  Fig. 
55.  The  bottle  A  contains  air.  B  contains  concentrated  sul- 
phuric acid,  C  contains  granulated  calcium  chloride,  D  is  care- 


788   EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XVI. 

fully  dried  and  contains  a  few  pieces  of  granulated  calcium 
chloride  and  air.  Pour  water  through  the  funnel-tube  into 
Ay  when  the  air  will  be  forced  through  B  and  C  and  into  D. 
But  in  passing  through  B  and  G  the  moisture  contained  in  it 


FIG.  55. 

will  be  removed,  and  the  air  which  enters  D  will  be  dry. 
After  A  has  once  been  filled  with  water,  empty  it  and  fill  it 
again,  letting  the  dried  air  pass  into  D.  This  operation  may 
be  repeated  indefinitely.  The  calcium  chloride  in  D  will  not 
grow  moist. 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XVI. 
PREPARATION  AND  PROPERTIES  OF  AMMONIA. 

Experiment  125. — To  a  little  ammonium  chloride  on  a 
watch-glass  add  a  few  drops  of  a  strong  solution  of  caustic 
soda,  and  notice  the  odor  of  the  gas  given  off.  Do  the  same 
thing  with  caustic  potash.  Mix  small  quantities  of  ammo- 
nium chloride  and  lime  in  a  mortar,  and  add  a  few  drops  of 
water. 

Experiment  126. — Mix  20  parts  iron  filings,  1  part  potas- 
sium nitrate,  and  1  part  solid  potassium  hydroxide,  and  heat 
the  mixture  in  a  test-tube.  Is  there  any  evidence  of  the 
formation  of  ammonia  ? 

Experiment  127. — Arrange  an  apparatus  as  shown  in  Fig. 
45.  In  the  flask  put  a  mixture  of  100.  grams  slaked  lime 
and  50  grams  ammonium  chloride.  Heat  on  a  sand-bath. 


AMMONIA. 


789 


After  the  air  is  driven  out,  the  gas  will  be  completely  ab- 
sorbed by  the  water  in  the  first  Wolff's 
flask  if  shaken  from  time  to  time.  Dis- 
connect the  delivery-tube  from  the  series 
of  Wolff's  flasks,  and  connect  with  an- 
other tube  bent  upward.  Collect  some  of 
the  gas  by  displacement  of  air,  placing  the 
vessel  with  the  mouth  doivmvard.  (Why  ?) 
The  arrangement  is  shown  in  Fig.  56.  The 
vessel  in  which  the  gas  is  collected  should 
be  dry,  as  water  absorbs  ammonia'  very 
readily.  Hence,  also,  it  cannot  be  collected 
over  water.  In  the  gas  collected  introduce 
a  burning  stick  or  taper.  Ammonia  does 
not  burn  in  air,  nor  does  it  support  com- 
bustion. In  working  with  the  gas  great  care 
must  be  taken  to  avoid  inhaling  it  in  any 
quantity.  After  enough  has  been  collected  in  cylinders  to 
'exhibit  the  chief  properties,  connect  the  delivery-tube  again 
with  the  series  of  Wolff's  flasks,  and  pass  the  gas  through  the 
water  as  long  as  it  is  evolved. 

AMMONIA  BURNS  IN  OXYGEN. 

Experiment  128. — Put  a  little  of  a  concentrated  solution 
of  ammonia  in  a  flask  placed  upon  a  tripod.  Heat  gently  and, 
from  a  gasometer,  pass  a  rapid  current  of  oxygen  through  a 
bent  tube  into  the  liquid.  Apply  a  light  to  the  mouth  of  the 
vessel,  when  the  ammonia  will  be  seen  to  burn. 


FIG.  56. 


AMMONIA  FORMS  AMMONIUM  SALTS  WITH  ACIDS. 

Experiment  129. — Put  100  cc.  dilute  ammonia  solution  in 
an  evaporating-dish.  Try  its  effect  on  red  litmus  paper. 
Slowly  add  dilute  hydrochloric  acid  until  the  alkaline  reaction 
is  destroyed  and  the  solution  is  neutral.  Evaporate  to  dry- 
ness  on  a  water-bath.  Compare  the  substance  thus  obtained 
with  sal-ammoniac,  or  ammonium  chloride.  Taste.  Heat 
on  a  piece  of  platinum  foil.*  Treat  with  a  caustic  alkali. 
Treat  with  a  little  concentrated  sulphuric  acid  in  dry  test- 
tubes.  Do  they  appear  to  be  identical  ?  Similarly  sulphuric 
acid  and  ammonia  yield  ammonium  sulphate  ;  nitric  acid  and 
ammonia  yield  ammonium  nitrate  ;  etc. 


790    EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XVL 

Experiment  13O. — Fill  a  dry  cylinder  with  ammonia  gas, 
and  another  of  the  same  size  with  hydrochloric  acid  gas.  Bring 
them  together  with  their  mouths  covered.  Quickly  remove 
the  covers,  when  a  dense  white  cloud  will  appear  in  and  about 
the  cylinders.  This  will  soon  settle  on  the  walls  of  the  vessels 
as  a  light  white  solid.  It  is  ammonium  chloride.  Thus,  from 
two  colorless  gases  we  get  a  solid  substance  by  an  act  of  chemi- 
cal combination.  Heat  is  evolved  in  the  act  of  combination. 

COMPOSITION  OF  AMMONIA. 

Experiment  131. — This  experiment  should  be  performed 
by  a  person  experienced  in  the  use  of  chemical  apparatus.  A 
glass-tube,  such  as  represented  in  Fig.  57,  provided  with  a 
glass  stop-cock  is  needed.  Fill  this  tube  with  chlorine  free 
from  air  over  a  saturated  solution  of  sodium  chloride.  After 

it  is  filled  let  it  stand  for  some 
time  mouth  downward  in  the 
solution  of  sodium  chloride  to 
let  the  liquid  drip  out  of  it. 
Close  the  stop-cock  and  re- 
move it  from  the  solution. 
Hold  the  tube  mouth  upward, 
and  pour  a  concentrated  solu- 
tion of  ammonia  into  the  fun- 
nel-like projection  above  the 
stop-cock,  put  in  the  glass 
stopper,  and  now  by  slightly 
opening  the  stop-cock  let  the 
ammonia  pass  drop  by  drop 
into  the  tube.  Reaction  be- 
tween the  chlorine  and  the 
ammonia  takes  place,  accom- 
panied by  a  marked  evolu- 
tion of  heat,  and  in  a  partly- 
darkened  room  light  is  seen. 
Great  care  must  be  taken  not 
to  admit  air  with  the  am- 
monia. After  nearly  all  the 
ammonia  has  passed  in  from 
the  funnel,  pour  into  the 
Fl°- 57>  funnel  about  two  thirds  as 

much    ammonia   as   has  already    been    used,  and    let    this 


PREPARATION  AND  PROPERTIES  OF  NITRIC  ACID.      791 

in  gradually.  Leave  the  stop-cock  closed,  and  fill  the 
funnel  with  dilute  sulphuric  acid. .  Fit  a  bent  tube  into  a 
cork,  fill  this  tube  with  dilute  sulphuric  acid :  put  the 
cork  in  the  funnel,  and  the  other  end  of  the  tube  in  a  small 
beaker  containing  dilute  sulphuric  acid,  and,  after  immersing 
the  long  tube  in  water  of  the  ordinary  temperature,  open  the 
stop-cock.  If  the  operation  has  been  carried  out  as  it  should 
be,  the  dilute  acid  will  flow  into  the  tube  until  it  is  two  thirds 
full,  and  will  then  stop.  The  residual  gas  is  nitrogen.  What 
evidence  in  regard  to  the  composition  of  ammonia  is  furnished 
by  this  experiment  ? 

The  arrangement  of  the  apparatus  in  the  last  stage  of  the 
experiment  is  shown  in  Fig.  57. 

PREPARATION  AND  PROPERTIES  OF  NITRIC  ACID. 

Experiment  132 — Arrange  an  apparatus  as  shown  in  Fig. 
58.     In  the  retort  put  20  grams  sodium  nitrate  (Chili  salt- 


FIG.  58. 

peter)  and  20  grams  concentrated  sulphuric  acid.  On  gently 
heating,  nitric  acid  will  distil  over,  and  be  condensed  in  the 
receiver.  After  the  acid  is  all  distilled  off,  remove  the  con- 
tents of  the  retort.  Kecrystallize  the  substance  from  water, 
and  compare  it  with  the  sodium  sulphate  obtained  in  the 
preparation  of  hydrochloric  acid.  (See  Experiment  74.)  In 
the  latter  stage  of  the  operation  the  vessels  become  filled  with 
a  reddish-brown  gas.  The  acid  which  is  collected  has  a  some- 
what yellowish  color. 


792    EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XVI. 

Experiment  133. — Mix  together  400  grams  concentrated 
sulphuric  acid  and  80  grams  ordinary  concentrated  nitric  acid. 
Pour  the  sulphuric  acid  into  the  nitric  acid.  Distil  the  mix- 
ture from  a  retort  arranged  as  in  the  preceding  experiment, 
taking  care  to  keep  the  neck  of  the  retort  cool  by  placing 
filter-paper  moistened  with  cold  water  on  it.  Use  the  acid 
thus  obtained  for  the  purpose  of  studying  the  properties  of 
pure  nitric  acid. 


NITKIC  ACID  GIVES  UP  OXYGEN  KEADILY,  AND  is  HENCE  A 
GOOD  OXIDIZING  AGENT. 

Experiment  134. — Pour  concentrated  nitric  acid  into  a 
wide  test-tube,  so  that  it  is  about  one-fourth  filled.  Heat 
the  end  of  a  stick  of  charcoal  of  proper  size,  and,  holding  the 
other  end  with  a  forceps,  introduce  the  heated  end  into  the 
acid.  It  will  continue  to  burn  with  a  bright  light,  even, 
though  it  is  placed  below  the  surface  of  the  liquid.  The 
action  is  oxidation.  The  charcoal  in  this  case  finds  the  oxy- 


Fio.  59. 

gen  in  the  acid  and  not  in  the  air.  Great  care  must  be 
taken  in  performing  this  experiment.  The  charcoal  should 
not  come  in  contact  with  the  sides  of  the  test-tube.  A  large 
beaker-glass  should  be  placed  beneath  the  test-tube,  so  that 
in  case  it  breaks  the  acid  will  be  caught  and  prevented  from 
doing  harm.  The  arrangement  of  the  apparatus  is  shown  in 
Fig.  59. 


NITRATES.  793 

The  gases  given  off  from  the  tube  are  offensive  and  poison- 
ous. Hence  this  experiment  as  well  as  all  others  with  nitric 
acid  should  be  carried  on  under  a  hood  in  which  the  draught 
is  good. 

Experiment  135. — Boil  a  little  strong  nitric  acid  in  a  test- 
tube  in  the  upper  part  of  which  some  horse-hair  has  been  in- 
troduced in  the  form  of  a  stopper.  The  horse-hair  will  take 
fire  and  burn,  and  leave  a  white  residue.  Hold  the  test-tube 
with  a  forceps  over  a  vessel  to  catch  the  contents  should  the 
tube  break. 

Experiment  136. — In  a  small  flask  put  a  few  pieces  of 
granulated  tin.  Pour  on  this  just  enough  strong  nitric  acid 
to  cover  it.  Heat  gently  over  a  small  flame.  Soon  action  will 
take  place.  Colored  gases  will  be  evolved,  the  tin  will  disap- 
pear, and  in  its  place  will  be  found  a  white  powder.  This 
consists  mostly  of  tin  and  oxygen.  (See  Experiment  13.) 

METALS  DISSOLVE  IN  NITRIC  ACID,  FORMING  NITRATES. 

Experiment  137. — Dissolve  a  few  pieces  of  copper-foil  in 
ordinary  commercial  nitric  acid  diluted  with  about  half  its 
volume  of  water.  The  operation  should  be  carried  on  in  a 
good-sized  flask  and  under  an  efficient  hood.  When  the  cop- 
per has  disappeared,  pour  the  blue  solution  into  an  evaporat- 
ing-dish,  and  evaporate  down  to  crystallization.  Compare  the 
substance  thus  obtained  with  copper  nitrate.  Heat  specimens 
of  each.  Treat  small  specimens  with  sulphuric  acid.  What 
evidence  have  you  that  the  two  substances  are  identical  ? 

NITRATES  ARE  DECOMPOSED  BY  HEAT. 

Experiment  138. — Heat  some  potassium  nitrate  in  a  test- 
tube.  .Introduce  a  piece  of  wood  with  a  spark  on  it.  Heat 
also  lead  nitrate,  copper  nitrate,  and  any  other  nitrates  which 
may  be  available.  What  difference  do  you  observe  between  the 
decomposition  of  potassium  nitrate  and  that  of  lead  nitrate  ? 

NITRATES  ARE  SOLUBLE  IN  WATER. 

Experiment  139. — Try  the  solubility  in  water  of  the  ni- 
trates used  in  the  last  experiment. 


794    EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XVI. 

NITRIC  ACID  is  REDUCED  TO  AMMONIA  BY  NASCENT 
HYDROGEN. 

Experiment  140 — In  a  good-sized  test-tube  treat  a  few 
pieces  of  granulated  zinc  with  dilute  sulphuric  acid.  What 
is  evolved?  Prove  it.  Now  add  drop  by  drop  dilute  nitric 
acid.  The  hydrogen  ceases  to  be  given  off.  Pour  the  con- 
,  tents  of  the  tube  into  an  evaporating-dish  and  evaporate  the 
liquid.  Put  the  residue  into  a  test-tube  and  add  caustic-soda 
solution,  when  the  smell  of  ammonia  will  be  noticed.  Try  the 
action  of  the  gas  on  red  litmus-paper.  Moisten  the  end  of  a 
glass  rod  with  a  little  hydrochloric  acid,  and  hold  it  in  the 
tube.  White  fumes  are  seen.  What  are  they  ?  Do  the  same 
with  nitric  acid.  What  are  the  fumes  in  this  case  ? 

NITROUS  ACID. 

Experiment  141. — Melt  25  grams  potassium  nitrate  in  a 
shallow  iron  plate  and  gradually  add  50  grams  metallic  lead 
cut  in  small  pieces.  Stir  them  together  as  thoroughly  as 
possible.  After  the  mass  is  cooled  down,  break  it  up  and 
treat  with  water  in  a  flask.  The  potassium  nitrite  will  dis- 
solve, while  the  lead  oxide  and  unused  lead  will  not  dissolve. 
Filter.  Add  a  little  sulphuric  acid  to  some  of  the  solution. 
A  colored  gas  will  be  given  off.  See  whether  a  solution  of 
potassium  nitrate  acts  in  the  same  way.  Treat  with  sulphu- 
ric acid  a  little  of  the  residue  left  after  heating  potassium 
nitrate  alone  in  a  test-tube  as  in  Experiment  138. 

NITROUS  OXIDE. 

Experiment  142. — In  a  retort  heat  10  to  15  grams  crystal- 
lized ammonium  nitrate  until  it  has  the  appearance  of  boiling. 
Do  not  heat  higher  than  is  necessary  to  secure  a  regular  evo- 
lution of  gas.  Connect  a  wide  rubber  tube  directly  with  the 
neck  of  the  retort  and  collect  the  evolved  gas  over  water,  as 
in  the  case  of  oxygen.  It  supports  combustion  almost  as  well 
as  pure  oxygen.  Try  experiments  with  wood,  a  candle,  and  a 
piece  of  phosphorus. 


OXIDES  OF  NITROGEN.  795 


NITRIC  OXIDE. 

Experiment  143. — Arrange  an  apparatus  as  shown  in  Fig. 
60.  In  the  flask  put  a  few  pieces  of  copper-foil.  Cover  this 
with  water.  Now  add  slowly,  waiting  each  time 
for  the  action  to  begin,  ordinary  concentrated 
nitric  acid.  When  enough  nitric  acid  has  been 
added  gas  will  be  evolved.  If  the  acid  is  added 
rapidly,  it  not  unfrequently  happens  that  the 
evolution  of  gas  takes  place  too  rapidly,  so  that 
the  liquid  is  forced  out  of  the  flask  through 
the  funnel-tube.  This  can  be  avoided  by  not 
being  in  a  hurry.  At  first  the  vessel  becomes 
filled  with  a  reddish-brown  gas,  but  soon  the 
gas  evolved  becomes  colorless.  Collect  over 

o 

water  two  or  three  vessels  full.      The  gas   col- 
lected is  principally  nitric  oxide,   NO,   though        FIG.  eo. 
it  is  frequently  mixed  with  a  considerable  quantity  of  nitrous 
oxide. 

Experiment  144. — Turn  one  of  the  vessels  containing  col- 
orless nitric  oxide  with  the  mouth  upward,  and  uncover  it. 
The  colored  gas  is  at  once  seen,  presenting  a  very  striking 
appearance.  Do  not  inhale  the  gas.  Perform  the  experi- 
ments with  nitric  oxide  where  there  is  a  good  draught. 

Experiment  145.— Pass  nitric  oxide  into  a  concentrated 
solution  of  ferrous  sulphate.  Afterwards  heat  the  solution 
and  collect  the  gas.  What  do  you  conclude  that  the  gas  is? 

NITROGEN"  TRIOXIDE.  . 

Experiment  146. — In  a  flask  fitted  with  a  safety-funnel  and 
a  delivery-tube  pour  nitric  acid  of  specific  gravity  1.30-1.35 
upon  coarsely  granulated  arsenious  oxide,  As203.  Heat  gently, 
and  conduct  the  gases  through  a  tube  surrounded  by  a  freez- 
ing mixture,  as  in  Experiment  110. 

NITROGEN  PEROXIDE. 

Experiment  147. — Admit  a  little  air  to  nitric  oxide  con- 
tained in  a  bell-jar  over  water,  and  let  the  ves'sel  stand. 
Almost  immediately  the  color  will  disappear,  showing  that  the 
nitrogen  peroxide  formed  is  decomposed.  Again  admit  air, 


796    EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XVII. 

and  let  the  vessel  stand.  The  same  changes  will  be  noticed  as 
in  the  first  instance.  If  oxygen  is  used  instead  of  air  the 
above  changes  can  be  repeated  over  and  over  again.  Devise 
an  experiment  for  the  purpose  of  determining  whether  the 
nitric  oxide  is  gradually  used  up  or  not. 


EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XVII. 

PHOSPHORUS. 

Experiment  148. — [This,  as  well  as  the  other  experiments 
with  phosphorus,  should  be  performed  only  by  an  experienced 
person.]  Arrange  an  apparatus  as  shown  in  Fig.  61.  The 
neck  of  the  retort  is  somewhat  drawn  out  and  bent  downward 
and  fitted  air-tight  by  means  of  a  cork  to  the  wide  glass  tube 
B.  Some  small  pieces  of  ordinary  phosphorus  are  now  care- 
fully slipped  into  the  retort — as  much 
as  is  obtained  by  cutting  up  two  sticks 
three  to  four  inches  long.  The  ap- 
paratus is  then  adjusted  as  shown  in 
the  figure,  so  that  the  end  of  the 
tube  B  dips  below  the  surface  of  the 
water  in  the  beaker  C.  The  whole  is 
then  allowed  to  stand  for  some  hours. 
The  oxygen  is  absorbed  from  the  air 
contained  in  the  vessel,  and  the  water 
rises  in  B.  Without  uncovering  the 
end  of  Bt  replace  the  water  in  C  by 
some  that  has  a  temperature  of  about 
50°.  Now  heat  the  retort  gradually, 
when  the  phosphorus  will  distil  over 
and  condense  in  C  in  the  molten  condition.  By  lowering  the 
heat  gradually  at  the  end  of  the  operation  it  can  finally  be 
stopped  without  danger  of  breaking. 

Experiment  149. — Dissolve  a  little  ordinary  phosphorus  in 
carbon  disulphide.  Pour  some  of  this  solution  upon  a  strip 
of  filter-paper,  and  let  this  hang  in  the  air  or  wave  it  gently 
in  the  air.  After  the  carbon  disulphide  has  evaporated  the 
phosphorus  will  take  fire. 

Experiment  150. — Bring  together  in  a  porcelain  crucible 
or  evaporating-dish  a  little  phosphorus  and  iodine.  It  will  be 
seen  that  simple  contact  is  sufficient  to  cause  the  two  sub- 


FIQ.  61. 


PHOSPHINE.  797 

stances  to  act  upon  each  other.     Direct  combination  takes 
place,  and  the  action  is  accompanied  by  light  and  heat. 

PHOSPHORUS  ABSTRACTS  OXYGEN  FROM  OTHER  SUBSTANCES. 

Experiment  151. — Add  a  little  of  a  solution  of  phosphorus 
in  carbon  disulphide  to  a  solution  of  copper  sulphate.  What 
change  takes  place  ? 

Experiment  152. — Put  a  few  pieces  of  ordinary  phosphorus 
in  a  glass  tube  and  seal  it.  Heat  gradually  to  300°.  Open 
the  tube  and  examine  the  product.  See  whether  it  takes  fire 
as  readily  as  ordinary  phosphorus  does ;  whether  it  dissolves 
in  carbon  disulphide ;  whether  it  melts  easily  when  put  in 
water  heated  to  between  45°  and  50°. 

PHOSPHINE. 

Experiment  153. — Arrange  an  apparatus  as  shown  in  Fig. 
<62.  In  the  small  flask  B  put  about  5  grams  caustic  potash  dis- 


FIG.  62. 

solved  in  10-15  cc.  water,  and  when  the  solution  is  cold  add  a 
few  small  pieces  of  phosphorus  the  size  of  a  pea.  Pass  hydrogen 
for  some  time  through  the  apparatus  from  the  generating- 
flask  A  until  all  the  air  is  displaced  ;  then  disconnect  at  D, 
leaving  the  rubber  tube,  closed  by  the  pinch-cock,  on  the 
tube  which  enters  the  flask.  Gently  heat  the  contents  of  the 
retort,  when  gradually  a  gas  will  be  evolved,  and  will  escape 
through  the  water  in  C.  As  each  bubble  comes  in  contact 


798    EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XVII. 

with  the  air  it  takes  fire,  and  the  products  of  combustion  ar 
range  themselves  in  rings,  which  become  larger  as  they  rise. 
They  are  extremely  beautiful,  particularly  if  the  air  of  the 
room  is  quiet.  Both  the  phosphorus  and  the  hydrogen  com- 
bine with  oxygen  in  the  act  of  burning.  Collect  some  of  the 
gas  in  a  tube  over  water,  and  then  place  the  tube  mouth  up- 
ward. What  difference  is  there  between  the  burning  of  the 
gas  under  these  circumstances,  and  that  noticed  when  the 
rings  are  formed?  Collect  another  tube  full  of  the  gas,  and 
let  this  stand  for  some  time.  Then  open  the  vessel  by  taking 
it  out  of  the  water.  Has  any  change  taken  place  in  the  gas  ? 


ARSENIC. 

Experiment  154 — Heat  a  small  piece  of  arsenic  on  charcoal 
in  the  flame  of  the  blow-pipe. 

ARSINE. 

Experiment  155 — Arrange  an  apparatus  as  shown  in  Fig. 
63.     Put  some  pure  granulated  zinc  in  the  flask  and  pour 


FIG.  68. 

dilute  sulphuric  acid  on  it.  The  calcium-chloride  tube  serves 
to  dry  the  gas.  When  the  air  is  all  out  of  the  vessel  and  the 
hydrogen  is  lighted,  add  slowly  a  little  of  a  solution  of  arsenic 
trioxide,  As203 ,  in  dilute  hydrochloric  acid.  The  appearance 
of  the  flame  will  soon  change.  It  will  become  paler,  with  a 
slightly  bluish  tint,  and  give  off  white  fumes.  (See  next  ex- 
periment.) 


MARSH  'S  TEST  FOR  ARSENIC— ANTIMONY,  ETC.     799 


MARSH'S  TEST  FOR  ARSENIC. 

Experiment  156. — Into  the  flame  of  the  burning  hydrogen 
and  arsine  produced  in  the  last  experiment  introduce  a  piece 
of  porcelain,  as  the  bottom  of  a  small  porcelain  dish  or  a  cru- 
cible, and  notice  the  appearance  of  the  spots.  Heat  by  means 
of  a  Buusen  burner  the  tube  through  which  the  gas  is  passing, 
which  should  be  of  hard  glass.  Just  beyond  the  heated 
place  there  will  be  deposited  a  thin  layer  of  metallic  arsenic, 
commonly  called  a  mirror  of  arsenic.  This  deposit  is  due  to 
the  direct  decomposition  of  the  arsine  into  arsenic  and  hydro- 
gen by  heat.  [Compare  ammonia,  phosphine,  and  arsine  with 
reference  to  their  stability.] 

ANTIMONY. 

Experiment  157. — Heat  a  small  piece  of  antimony  on  char- 
coal in  the  blow-pipe  flame.  Try  the  action  of  dilute  and  of 
concentrated  hydrochloric  acid,  of  dilute  and  of  concentrated 
nitric  acid,  and  of  a  mixture  of  the  two  acids  on  a  small  piece 
of  antimony. 

STIBINE. 

Experiment  158. — Stibine  is  made  by  the  same  method  as 
that  used  in  making  arsine.  Make  some,  using  a  solution  of 
tartar  emetic.  Introduce  a  piece  of  porcelain  in  the  flame, 
and  afterwards  heat  the  tube  through  which  the  gas  is  passing. 
Compare  the  antimony  spots  with  the  arsenic  spots.  Color? 
Volatility?  Conduct  towards  a  solution  of  sodium  hypochlo- 
rite  or  hypobromite? 

BISMUTH. 

Experiment  159. — Heat  a  piece  of  bismuth  on  charcoal  in 
the  blow-pipe  flame.  See  how  it  conducts  itself  towards  hy- 
drochloric acid;  towards  nitric  acid.  If  a  solution  is  obtained 
in  either  case,  add  water  to  it.  Explain  what  takes  place. 

PHOSPHORUS  TRICHLORIDE. 

The  experiments  with  the  chlorides  of  phosphorus  must  be 
carried  on  under  a  hood  or  out-of-doors. 

Experiment  160. — Arrange  an  apparatus  as  shown  in  Fig. 
64.  The  tube  A  is  arranged  so  that  it  can  be  raised  or  low- 
ered in  the  retort.  Put  50  to  100  grams  ordinary  phosphorus 
in  the  retort,  taking  precautions  to  prevent  it  from  taking  fire 


800     EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XVIL 
during  the  operation.     This  is  best  accomplished  by  fitting 


FIG.  64. 

corks  in  both  openings  of  the  retort;  placing  the  retort  in  a 
vessel  of  cold  water;  removing  the  cork  from  B,  throwing  in 
a  piece  of  phosphorus,  and  quickly  putting  the  cork  in.  The 
pieces  must  not  be  put  in  in  too  rapid  succession.  After  all 
the  phosphorus  is  in  the  retort,  adjust  the  apparatus  as  repre- 
sented, placing  the  receiver  D  in  a  dish  of  cold  water.  Now 
connect  by  means  of  the  rubber  tube  E  with  an  apparatus 
furnishing  chlorine,  dried  by  means  of  concentrated  sulphuric 
acid  and  calcium  chloride.  As  soon  as  the  chlorine  comes  in 
contact  with  the  phosphorus  action  begins,  and  the  product, 
which  is  phosphorus  trichloride,  distils  over  into  the  receiver. 
If  the  action  is  taking  place  too  rapidly,  the  inside  of  the  re- 
tort will  become  covered  with  a  coating  of  red  phosphorus. 
In  this  case  raise  the  tube  A  a  little  and  the  red  coating  will 
gradually  disappear.  If  the  tube  is  raised  too  high,  not  enough 
heat  is  generated,  and  the  trichloride  in  the  retort  is  converted 
into  the  pentachloride,  which  is  deposited  as  a  white  coating. 
By  raising  and  lowering  the  tube  ac- 
cording to  the  indications,  the  retort 
can  be  kept  clear,  and  all  the  phos- 
phorus converted  into  the  trichloride. 
This  manipulation  of  the  tube  is  much 
facilitated  by  fitting  into  the  cork  a 
somewhat  larger  tube,  through  which 
the  smaller  one  can  pass  easily;  letting 
this  project  about  an  inch  and  an  half 
above  the  cork  and  passing  over  it  a 
piece  of  rubber  tubing  of  such  size 
that  while  the  smaller  tube  moves  through  it  readily,  the  two 
form  a  gas-tight  joint.  This  is  shown  in  Fig.  65.  After  the 


FIG.  65. 


PHOSPHORUS  PENTACHLORIDE. 


801 


operation  is  finished,  pour  the  liquid  from  the  receiver  into  a 
clean  dry  flask,  and  distil  on  a  water-bath.  Try  the  action  of 
a  little  of  the  compound  on  water. 

PHOSPHORUS  PENTACHLORIDE. 

Experiment  161 — Put  the  trichloride  of  phosphorus  ob- 
tained in  the  last  experiment  in  a  wide-mouthed 
bottle  surrounded  by  cold  water.  Through  a 
wide  glass  tube  pass  dry  chlorine  upon  the 
surface  of  the  liquid,  and  as  the  action  ad- 
vances, and  a  solid  begins  to  make  its  appear- 
ance, stir  the  contents  of  the  bottle.  Con- 
tinue the  passage  of  the  chlorine  until  the 
product  is  a  perfectly  dry  solid.  The  arrange- 
ment of  the  bottle  containing  the  trichloride, 
and  that  of  the  delivery-tube,  is  shown  in  Fig. 
66.  The  bottle  is  put  in  a  larger  vessel  con- 
taining cold  water,  which  is  renewed  from 
time  to  time  during  the  process. 

Try  the  action  of  a   little   phosphorus  pentachloride  on 
water.     In  a  large  dry  flask  heat  a  little  of  the  pentachloride. 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XVIII. 

PHOSPHORIC  ACID. 
Experiment  162. — In  a  flask  connected  with  an  inverted 


FIG.  66. 


FIG  67. 


condenser,  as  shown  in  Fig.  67,  boil  10  to  15  grams  of  ordinary 
phosphorus  with  250  cc.  ordinary  commercial  nitric  acid.     If 


802    EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XVIIL 

necessary,  add  more  acid  after  a  time.  Boil  gently  until  the 
phosphorus  disappears.  Evaporate  the  solution  to  complete 
dryness,  so  as  to  get  rid  of  all  the  nitric  acid.  Dissolve  a  lit- 
tle of  the  product  in  water,  and  add  a  few  drops  of  the  solu- 
tion to  a  dilute  solution  of  silver  nitrate.  What  effect  is  pro- 
duced? Heat  some  of  the  product  gently  in  a  porcelain 
crucible,  and  from  time  to  time  take  out  a  little,  dissolve  it  in. 
water,  and  try  its  action  on  silver  nitrate. 

Experiment  163. — Try  the  action  of  ordinary  sodium  phos- 
phate on  silver  nitrate.  Heat  a  little  of  the  salt  in  a  porcelain 
crucible  to  redness.  After  cooling,  try  the  action  of  the  salt 
left  in  the  crucible  on  silver  nitrate. 

ARSENIC  ACID. 

Experiment  164. — Pass  chlorine  into  water  containing  ar- 
senic trioxide  in  suspension,  until  the  oxide  is  dissolved. 
Evaporate  to  crystallization.  Into  a  dilute  solution  of  the 
product  thus  obtained,  to  which  some  hydrochloric  acid  is 
added,  pass  hydrogen  sulphide.  Explain  the  changes. 

KEDUCTION  OF  ARSENIC  TRIOXIDE. 

Experiment  165. — In  the  bottom  of  a  dry  tube  of  hard  glass 
of  the  form  represented  in  Fig.  68  put  a  minute  piece  of  ar- 
senic trioxide,  and  just  above  it  a  small  bit  of  charcoal. 
Heat  gently.     Explain  the  change. 

SULPHIDES  OF  ARSENIC. 

Experiment  166. — Pass  hydrogen  sulphide  into  a  di- 
lute solution  of  arsenic  trioxide  in  hydrochloric  acid. — 
Filter  off  the  precipitate,  and  try  the  action  of  ammoni- 
um sulphide  on  some  of  it. 

SULPHIDES  OF  ANTIMONY. 

Experiment  167. — Pass  hydrogen  sulphide  into  a  so- 
FIO.  68.  ju^on  Of  antimonio  acid  made  by  treating  antimony 
with  aqua  regia  and  diluting  with  water.  Pass  hydrogen  sul- 
phide into  a  solution  of  antimony  trichloride  made  by  dis- 
solving stibnite  or  antimony  trisulphide  in  hydrochloric  acid. 
Try  the  action  of  ammonium  sulphide  on  the  precipitates 
after  filtering. 

OXYCHLORIDES   OF   ANTIMONY. 

Experiment  168. — Treat  a  solution  of  antimony  trichloride 
with  water. 


BASIC  NITRATES  OF  BISMUTH— CARBON.  803 

BASIC   NITRATES   OF   BlSMUTH. 

Experiment  169. — Dissolve  a  little  bismuth  in  nitric  acid 
and  evaporate.  Add  water. 

BORON. 

Experiment  170. — Make  a  hot  solution  of  30  grams  crystal- 
lized borax  in  120  cc.  water.  Add  slowly  10  grams  concen- 
trated sulphuric  acid.  On  cooling,  the  boric  acid  will  crystal- 
lize out.  What  evidence  have  you  that  the  substance  which 
crystallizes  out  of  the  solution  is  not  borax?  Try  the  solu- 
bility in  alcohol  of  specimens  of  each.  Is  there  any  difference? 
Treat  a  few  crystals  of  borax  with  about  10  cc.  alcohol  ;  pour 
off  the  alcohol  and  set  fire  to  it.  Treat  a  few  crystals  of  the 
boric  acid  in  the  same  way.  What  difference  do  you  observe  ? 
Distil  an  aqueous  solution  of  boric  acid,  and  determine 
whether  any  of  the  acid  passes  over  with  the  water  vapor. 

EXPERIMENTS  TO  ACCOMPANY   CHAPTER  XIX. 

CARBON. — BONE-BLACK  FILTERS. 

Experiment  171. — Make  a  filter  of  bone-black  by  fitting  a 
paper  filter  into  a  funnel  12  to  15  mm.  (5  to  6  inches)  in  di- 
ameter at  its  mouth.  Half  fill  this  with  bone-black.  Pour  a 
dilute  solution  of  indigo  through  the  filter.  If  the  conditions 
are  right  the  solution  will  pass  through  colorless.  Do  the 
same  thing  with  a  dilute  solution  of  litmus.  If  the  color  is 
not  completely  removed  by  one  filtration,  heat  and  filter 
again.  The  color  can  also  be  removed  from  solutions  by  put- 
ting some  bone-black  into  them  and  boiling  for  a  time.  Try 
this  with  half  a  liter  each  of  the  litmus  and  indigo  solutions 
used  in  the  first  part  of  the  experiment.  Use  about  4  to  5 
grams  bone-black  in  each  case.  Shake  the  solution  frequently 
while  heating. 

CHARCOAL  ABSORBS  GASES. 

Experiment  172. — Collect  over  mercury  in  glass  tubes  some 
ammonia  gas,  and  some  carbon  dioxide.  Introduce  into  each 
a  piece  of  charcoal,  which  has  been  heated  in  a  Bunsen- 
burner  flame  in  order  to  drive  out  gases  which  may  be  con- 
tained in  the  pores.  •*•  » 

CARBON  COMBINES  WITH  OXYGEN  TO  FORM  CARBON  DIOXIDE. 

Experiment  173 — Put  a  small  piece  of  charcoal  in  a  piece  of 

hard-glass  tube.     Heat  the  tube,  and  pass  oxygen  through  it. 


804   EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XIX. 


Pass  the  gases  into  clear  lime-water.     Arrange  the  apparatus 
as  shown  in  Fig.  69. 


FIG.  69. 

A  is  a  large  bottle  containing  oxygen  ;  B  is  a  cylinder  con- 
taining sulphuric  acid;  0  is  a  U-tube  containing  calcium 
chloride ;  D  is  the  hard-glass  tube  containing  the  charcoal; 
E  is  the  cylinder  with  clear  lime-water.  Explain  all  that 
takes  place. 

CARBON  REDUCES  SOME  OXIDES  WHEN  HEATED  WITH  THEM. 

Experiment  174 — Mix  together  two  or  three  grams  pow- 
dered copper  oxide,  CuO,  and  about  one  tenth  its  weight  of 

powdered  charcoal ;  heat  in  a  hard- 
glass  tube,  as  shown  in  Fig.  70,  or, 
still  better,  use  an  arsenic-tube. 

Pass  the  gas  which  is  given  off 
into  lime-water  contained  in  a  test- 
tube.  Is  it  carbon  dioxide  ?  What 
evidence  have  you  that  oxygen  has 
been  extracted  from  the  copper 
oxide  ?  Compare  the  substance  left 
in  the  tube  with  metallic  copper. 
Treat  both  with  nitric  acid,  with 
sulphuric  acid. 

Experiment  175. — Eepeat  Experiment  165  with  somewhat 
larger  quantities  of  the  substances,  and  examine  the  gas 
given  off. 

HYDROCARBONS. 

Experiment  176. — Make  marsh-gas  by  heating  in  a  retort  a 


FIG.  70. 


CARBON  DIOXIDE. 


805 


mixture  of  20  grams  sodium  acetate,  20  grams  potassium  hy- 
droxide, and  30  grams  slaked  lime.  Collect  some  of  the  gas 
over  water.  Is  it  a  combustible  gas  ? 

Experiment  177. — Make  ethylene  as  follows  :  In  a  flask  of 
2  to  3  liters  capacity  put  a  mixture  of  25  grams  alcohol  and 
150  grams  ordinary  concentrated  sulphuric  acid.  Heat  to 
160°  to  170°,  and  add  gradually  through  a  funnel  tube  about 
500  cc.  of  a  mixture  of  1  part  of  alcohol  and  2  parts  of  con- 
centrated sulphuric  acid.  Pass  the  gas  through  three  wash- 
bottles  containing,  in  order,  concentrated  sulphuric  acid, 
caustic  soda,  and  concentrated  sulphuric  acid.  Collect  some 
of  the  gas  over  water.  Is  it  combustible  ? 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XX. 
CARBON  DIOXIDE   is  FORMED  WHEN  A  CARBONATE  is 

TREATED    WITH   AN    ACID. 

Experiment  178. — In  test-tubes  add  successively  dilute  hy- 
drochloric, sulphuric,  nitric,  and  acetic  acids  to  a  little  sodium 
carbonate.  In  each  case  pass  the  gas  given  off  through  lime- 
water,  and  insert  a  burning  stick  in  the  upper  part  of  each 
tube. — Perform  the  same  experiments  with  small  pieces  of 
marble. 

PREPARATION  AND  PROPERTIES  OF  CARBON  DIOXIDE. 
Experiment  179. — Arrange  an  apparatus  as  shown  in  Fig. 
71.     In  the  flask  put  some  pieces  of  marble  or  limestone, 
and   pour   ordinary  hydrochloric  acid 
on  it.     The  gas  should  be  collected  by 
displacement  of  air,  the  vessel  being 
placed  with  the  mouth  upward.     Col- 
lect several  cylinders  or  bottles  full  of 
the   gas.     Into   one   introduce  succes- 
sively a  lighted  candle,  a  burning  stick, 
a  bit   of   burning   phosphorus.       Into 
another,  if  convenient,  put  a  live  mouse. 
"With   another   proceed  as   if   pouring 
water  from  it.     Pour  the  invisible  gas 
upon  the  flame  of  a  burning  candle. 
Pour  some  of  the  gas  from  .  one  vessel 
to  another,  and  show  that  it  has  been  ^ 
transferred.      Balance   a   beaker   on  a  ~~  "~^™— 

good-sized  pair  of  scales,  and  pour  car- 
bon dioxide  into  it.     If  the  balance  is  at  all  sensitive,  the  pan 
on  which  the  beaker  is  placed  will  sink. 


806    EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XX. 

CARBON  DIOXIDE  is  GIVEN  OFF  FROM  THE  LUNGS. 

Experiment  180. — Force  the  gases  from  the  lungs  through 

some  lime-water  by  means  of  an 
apparatus  arranged  as  shown  in 
Fig.  72. 

FORMATION  OF  CARBONATES. 

Experiment  181. — Pass  car- 
bon dioxide  into  a  solution  of  po- 
tassium hydroxide  to  saturation. 
Determine  whether  a  carboiiate  is 
in  solution  or  not. 

Experiment  182. — Pass  car- 
bon dioxide  into  50  to  100  cc.  clear 
lime-water.  Filter  off  the  white 

insoluble  substance.     Try  the  action   of  a  little  acid  on  it. 

What  evidence  have  you  that  it 'is  a  carbonate  ? 

Experiment  183. — Pass  carbon  dioxide  first  through  a  little 

water  to  wash  it,  and  then  into  50  to  100  cc.  clear  lime-water. 

Continue  to  pass  the  gas  for  some  time  after  the  precipitate 

is  formed.      The   precipitate  dissolves.     Heat   the   solution. 

What  happens  ?     Explain  these  reactions. 

PREPARATION  AND  PROPERTIES  OF  CARBON  MONOXIDE. 

Experiment  184 — Put  10  grams  crystallized  oxalic  acid  and 
50  to  60  grams  concentrated  sulphuric  acid  in  an  appropriate 
flask.  Connect  with  two  Wolff's  flasks  containing  a  solution  of 
caustic  soda,  so  that  the  gas  evolved  will  bubble  through  the 
solution.  Heat  gently.  Collect  some  of  the  gas  over  water. 
Set  fire  to  some,  and  notice  the  characteristic  blue  flame.  If 
convenient  put  a  live  mouse  in  a  vessel  containing  a  mixture 
of  about  equal  parts  of  carbon  monoxide  and  air.  It  will  die 
unless  taken  out. 

CARBON  MONOXIDE  is  A  GOOD  KEDUCING  AGENT. 

Experiment  185. — Pass  carbon  monoxide  over  some  heated 
copper  oxide  contained  in  a  hard-glass  tube.  Is  the  oxide  re- 
duced? How  do  you  know?  Is  carbon  dioxide  formed? 
What  evidence  have  you  ?  Was  the  carbon  monoxide  used 
free  of  carbon  dioxide?  If  not,  what  evidence  have  you  that 
carbon  dioxide  is  formed  in  this  experiment  ? 

Experiment  186. — Pass  carbon  dioxide  over  heated  charcoal 
in  a  hard-glass  tube.  What  is  formed  ? 


COAL-GAS,  ETC.  807 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXI. 
COAL-GAS. 

Experiment  187. — Heat  some  bituminous  coal  in  a  retort 
and  collect  over  water  the  gases  given  off.  Are  these  gases 
combustible  ? 

OXYGEN  BURNS  IN  AN  ATMOSPHERE  OF  A  COMBUSTIBLE  GAS. 

Experiment  188. — Break  off  the  neck  of  a  good-sized  re- 
tort ;  fit  a  perforated  cork  to  the  small  end  ;  pass  a  piece  of 
glass  tube  through  the  cork,  and  connect  by  means 'of  rubber 
hose  with  an  outlet  for  coal-gas.  Fix  the  apparatus  in  position, 


FIG.  73. 

as  shown  in  Fig.  73.  Turn  the  gas  on,  and  when  the  air  is 
driven  out  of  the  retort-neck,  light  the  gas.  The  neck  is  now 
filled  with  illuminating  gas,  and  the  gas  is  burning  at  the  mouth 
of  the  vessel.  If  now  a  platinum  jet  from  which  oxygen  is 
issuing  is  passed  up  into  the  gas  the  oxygen  will  take  fire,  and  a 
flame  will  appear  where  the  oxygen  escapes  from  the  jet. 
The  oxygen  burns  in  the  atmosphere  of  coal-gas. 

KINDLING  TEMPERATURE  OF  GASES. 

Experiment  189 — Light  a  Bunsen  burner.  Bring  down 
upon  the  flame  a  piece  of  brass  or  iron  wire-gauze.  There  is 
no  flame  above  the  gauze.  That  the  gas  passes  through  un- 
burned  can  be  shown  by  applying  a  light  just  above  the  outlet 
of  the  burner  and  above  the  gauze.  The  gas  will  take  fire  and 
burn.  By  simply  passing  through  the  thin  wire-gauze,  then,  the 
gas  is  cooled  down  below  its  burning  temperature,  and  does  not 
burn  unless  it  is  heated  up  again.  Turn  on  a  Bunsen  burner. 


808    EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXL 


Do  not  light  the  gas.  Hold  a  piece  of  wire-gauze  about  one 
and  a  half  to  two  inches  above  the  outlet.  Apply  a  lighted 
match  above  the  gauze,  when  the  gas  will  burn 
above  the  gauze,  but  not  below  it.  Here  again 
the  heat  necessary  to  raise  the  temperature  of  the 
gas  to  the  burning  temperature  cannot  be  com- 
municated through  the  gauze.  If  in  either  of  the 
above-described  experiments  the  gauze  is  held  in 
position  for  a  time,  it  will  probably  become  so 
highly  heated  that  the  gas  on  the  side  where  there 
is  no  flame  will  be  raised  to  the  burning  tempera- 
ture. The  instant  that  point  is  reached  the 
flame  becomes  continuous. 

THE  BLOW-PIPE  AND  ITS  USES. 

FIG  ?4  The  blow-pipe  used  in  chemical  laboratories  is 

constructed  as  shown  in  Fig.  74. 

When  used  with  the  Bunsen  burner  it  is  best  to  slip  into 
the  burner  a  brass  tube  ending  above 
in  a  narrow  slit-like  opening,  as  shown  in 
Fig.  75.  The  tube  referred  to,  marked 


FIG.  76. 

a  in  the  figure,  reaches  to  the  bottom  of 
the  burner,  and  thus  cuts  off  the  supply 
,  of  air  which  usually  enters  the  holes  at 
FIG.  75.  the  base.     The  gas  is  now  lighted,  and 

the  current  so  regulated  that  there  is  a  small  flame  about  1^- 
to  2  inches  long.  The  tip  of  the  blow-pipe  is  placed  on  the 
slit  of  the  burner  in  the  flame,  as  shown  in  Fig.  76.  By  blow- 
ing regularly  and  not  violently  through  the  pipe  the  flame  is 
forced  down  in  the  same  direction  as  the  end-piece  of  the 
blow-pipe,  and  the  slant  of  the  burner-slit.  Under  proper 


THE  BLOW-PIPE  AND  ITS  USES,  809 

conditions  the  flame  separates  sharply  into  a  central  blue  part 
and  an  outer  part  of  another  color.  The  direction  and  lines  of 
division  of  the  flame  are  indicated  in  Fig.  76.  The  outer 
part  of  the  flame  marked  o  is  the  oxidizing  flame  ;  the  part 
marked  r  is  the  reducing  flame. 

Experiment  190. — Select  a  piece  of  charcoal  about  4 
inches  long  by  1  inch  wide  and  1  inch  thick,  with  one  surface 
plane.*  Near  the  end  of  the  plane  surface  make  a  cavity  by 
pressing  the  edge  of  a  small  thin  coin  against  it,  and  turning 
it  completely  round  a  few  times.  Mix  together  equal  small 
quantities  of  dry  sodium  carbonate  -and  lead  oxide.  Put  a 
little  of  the  mixture  in  the  cavity  in  the  charcoal,  and  heat  it 
in  the  reducing  flame  produced  by  the  blow-pipe.  In  a  short 
time  globules  of  metallic  lead  will  be  seen  in  the  molten  mass. 
After  cooling,  scrape  the  solidified  substance  out  of  the  cavity 
in  the  charcoal.  Put  it  in  a  small  mortar,  treat  it  with  a  little 
water,  and,  after  breaking  it  up  and  allowing  as  much  as  pos- 
sible to  dissolve,  pick  out  the  metallic  beads.  Is  it  malleable 
or  brittle?  Is  metallic  lead  malleable  or  brittle?  Is  it  dis- 
solved by  hydrochloric  acid  ?  Is  lead  soluble  in  hydrochloric 
acid  ?  Is  it  soluble  in  nitric  acid  ?  Is  lead  soluble  in  nitric 
acid  ?  The  action  of  the  acids  can  be  tried  by  putting  the 
bead  on  a  small  dry  watch-glass  and  adding  a  few  drops  of  the 
acid.  Does  the  substance  act  like  lead?  What  has  become 
of  the  oxygen  with  which  the  lead  was  combined  in  the  oxide  ? 
Is  there  any  special  advantage  in  having  a  support  of  charcoal 
for  this  experiment? 

Experiment  191. — Heat  a  small  piece  of  metallic  lead  on 
charcoal  in  the  oxidizing  blow-pipe  flame.  Notice  the  forma- 
tion of  the  oxide,  which  forms  a  coating  or  film  on  the  char- 
coal in  the  neighborhood  of  the  metal.  Is  there  any  analogy 
between  this  process  and  the  burning  of  hydrogen  ?  In  what 
does  the  analogy  consist?  What  differences  are  there  between 
the  two  processes? 

Experiment  192 — Repeat  the  experiments  with  arsenic, 
antimony,  and  bismuth.  Notice  the  colors  of  the  films  formed 
on  the  charcoal. 

Experiment  193.— Melt  into  a  bit  of  glass  tubing  a  piece  of 
platinum  wire  8  to  10  mm.  (3  to  4  inches  long)  and  bend  the 
•end  so  as  to  form  a  small  loop,  as  shown  in  Pig.  77.  Heat  the 

*  Pieces  of  charcoal  prepared  for  blow -pipe  work  can  be  bought  from 
dealers  in  chemical  apparatus,  at  small  cost. 


810    EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXll. 

loop  in  the  flame  of  a  Bunsen  burner,  and  then  dip  it  into 
some  sodium-ammonium  phosphate  (microcosmic  salt).  Heat 
in  the  oxidizing  flame  of  the  blow-pipe  until  a  clear  glass  bead 
is  formed  in  the  loop.  What  changes  have  taken  place  ?  and 
what  is  the  clear  glass  ?  Bring  a  minute  particle  of  a  man- 


FIG.  77. 

ganese  compound  in  contact  with  the  bead,  and  heat  again. 
What  change  takes  place  ?  Try  the  same  experiment,  using 
successively  a  cobalt  compound,  a  copper  compound,  and  an 
iron  compound.  Now,  instead  of  using  microcosmic  salt,  use 
borax.  Explain  the  changes  in  all  the  above-described  experi- 
ments. 

CYANOGEN. 

Experiment  194. — Make  potassium  cyanide  by  heating  po- 
tassium ferrocyanide  in  an  iron  crucible. 

Experiment  195. — Make  cyanogen  by  heating  mercuric  cy- 
anide. Cyanogen  is  poisonous.  Burn  some  of  the  gas. 

Experiment  196. — Make  potassium  cyanate  from  some  of 
the  cyanide  obtained  in  Experiment  194.  This  is  done  by 
melting  it  in  an  iron  crucible,  and,  while  the  mass  is  liquid, 
adding  about  four  times  its  weight  of  red  lead,  stirring  during 
the  operation.  After  this  the  crucible  should  again  be 
put  in  the  furnace  for  a  little  while,  the  metallic  lead  allowed 
to  settle,  and  the  contents  poured  out  on  a  smooth  stone. 
Break  this  up,  and  extract  the  cyanate  with  alcohol. 

EXPERIMENTS  TO  ACCOMPANY   CHAPTER  XXII. 
SILICON. 

Experiment  197. — Prepare  sodium  fluosilicate  as  directed 
in  the  next  experiment.  Mix  3  parts  of  the  dry  salt  with  1 
part  of  sodium  cut  in  pieces.  Throw  this  mixture  all  at  once 
into  a  Hessian  crucible  heated  to  bright-red  heat  in  a  furnace. 
Add  immediately  9  parts  granulated  zinc,  and  a  layer  of  sodium 
chloride  previously  heated  to  drive  off  water.  The  crucible  is 
then  covered,  and  the  fire  allowed  to  burn  down.  After  cool- 
ing, the  regulus  of  zinc  containing  the  silicon  is  separated  from 
the  slag,  washed  with  water,  and  treated  with  hydrochloric 


SILICON  TETRAFLUORIDE  AND  FLUOSILICIC  ACID.     811 

acid.  The  zinc  dissolves  and  leaves  the  silicon.  This  is  again 
washed  with  water  and  then  heated  with  nitric  acid,  and 
washed  with  water,  when  crystals  of  silicon,  sometimes  of 
great  beauty,  are  obtained.  Try  the  effect  of  heating  a  little 
of  the  silicon  in  the  air.  Try  the  action  of  acids  and  of 
alkalies  upon  it. 

SILICON  TETKAFLUORIDE  AND  FLUOSILICIC  ACID. 

Experiment  198. — Arrange  an  apparatus  as  shown  in  Fig. 
78.     A  is  a  bottle  of  about  2  liters  capacity,  such  as  are  com- 

Ji 


FIG.  78. 

monly  used  for  transporting  acids.  This  is  about  two-thirds 
filled  with  alternating  layers  of  sand  and  powdered  fluor-spar, 
moistened  with  concentrated  sulphuric  acid.  The  bottle  is 
put  in  the  deep  sand-bath  B,  and  connected  by  means  of  a 
wide  glass  tube  with  the  funnel  C,  which  dips  just  below  the 
surface  of  the  water  in  the  large  evaporating-dish  D.  The 
sand-bath  is  now  gently  heated,  when  silicon  tetrafluoride 
passes  over.  Coming  in  contact  with  water,  it  is  decomposed, 
silicic  acid  being  deposited  -and  fluosilicic  acid  passing  into 
solution.  In  order  to  prevent  clogging,  the  gelatinous  silicic 
acid  is  from  time  to  time  removed  from  the  mouth  of  the 
funnel  by  means  of  a  bent-glass  rod.  After  the  action  is  com- 
plete, filter  the  solution.  Take  out  one  quarter,  and  to  the 


812     EXPERIMENTS  TO  ACCOMPANY  CHAPTER 

rest  slowly  add  a  solution  of  sodium  carbonate  until  the  whole 
just  begins  to  show  an  alkaline  reaction;  now  add  the  other 
quarter  of  the  acid,  and  filter.  Explain  all  the  reactions. 
Heat  a  little  of  the  dried  salt  in  a  covered  platinum  crucible. 
What  change  takes  place?  What  evidence  have  you  that  the' 
change  has  taken  place  ?  To  a  little  of  the  salt  in  water  add 
a  solution  of  potassium  hydroxide.  What  change  takes  place  ? 
Dry  the  silicic  acid  formed  in  the  first  part  of  the  experiment 
by  decomposition  of  the  silicon  tetrafluoride. 

SILICIC  ACID. 

Experiment  199. — Boil  some  of  the  silicic  acid  obtained  in 
the  last  experiment  with  sodium  hydroxide.  Treat  some  of 
the  solution  with  hydrochloric  acid  ;  with  ammonium  chloride. 

Experiment  200. — Add  some  fine  sand  to  about  four  times 
its  weight  of  a  molten  mixture  of  potassium  and  sodium  car- 
bonates, heated  in  a  platinum  crucible  in  the  flame  of  the 
blast-lamp.  Continue  the  heating  until  no  more  sand  is  dis- 
solved. Pour  the  molten  mass  out  on  a  stone,  and  when 
cooled  break  it  up  and  treat  it  with  water. 

Experiment  201. — Treat  a  little  of  the  solution  containing 
sodium  and  potassium  silicates,  prepared  in  the  last  experi- 
ment, with  a  little  sulphuric  or  hydrochloric  acid.  A  gelati- 
nous substance  will  be  precipitated.  This  is  silicic  acid.  Some 
of  the  acid  remains  in  solution.  By  evaporating  the  solution 
to  dryness  and  heating  for  a  time  on  the  water-bath,  all  the 
silicic  acid  is  rendered  insoluble. 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXIV. 
CHLORIDES,  BROMIDES,  AND  IODIDES. 

Experiment  202. — Dissolve  a  small  crystal  of  silver  nitrate 
in  pure  water.  Add  to  a  small  quantity  of  this  solution  in  a 
test-tube  a  few  drops  of  dilute  hydrochloric  acid.  The  white 
substance  thus  precipitated  is  silver  chloride,  AgCl.  To  an- 
other small  portion  of  the  solution  add  a  few  drops  of  a  dilute 
solution  of  common  salt,  or  sodium  chloride,  NaCl.  The 
white  substance  produced  in  this  case  is  also  silver  chloride. 
Add  ammonia  to  each  tube.  If  sufficient  is  added  the 
precipitates  will  dissolve.  On  adding  enough  hydrochloric 
acid  to  these  solutions  to  combine  with  all  the  ammonia  the 


HYDROXIDES.  813 

silver  chloride  is  again  thrown  down.  On  standing  exposed  to 
the  light  both  precipitates  change  color,  becoming  finally  dark 
violet.  The  reactions  involved  in  the  above  experiments  are 
these  :  In  the  first  place,  when  hydrochloric  acid  is  added  to 
silver  nitrate  this  reaction  takes  place  : 

AgNO,  +  HC1  =  AgCl  +  HN03. 

When  sodium  chloride  is  added  this  reaction  takes  place  : 
AgN03  +  NaCl  =  AgCl  -f  NaNO,. 

In  the  first  reaction  nitric  acid  is  set  free  ;  in  the  second, 
the  sodium  and  silver  exchange  places.  In  addition  to  the 
insoluble  silver  chloride,  there  is  formed  at  the  same  time  the 
soluble  salt,  sodium  nitrate.  On  adding  ammonia  the  silver 
chloride  forms  with  it  a  compound  which  is  soluble  in  water ; 
and  on  adding  an  acid,  the  ammonia  combines  with  it,  leav- 
ing the  silver  chloride  uncombined  and  therefore  insoluble. 

Extensive  use  is  made  of  insoluble  compounds  for  the  pur- 
pose of  detecting  substances  in  analysis.  The  only  insoluble 
chlorides  are  those  of  silver,  lead,  and  mercury.  *  If,  there- 
fore, on  adding  hydrochloric  acid  or  a  soluble  chloride  to  a 
solution,  a  precipitate  is  formed,  the  conclusion  is  justified 
that  one  or  more  of  the  three  metals — silver,  lead,  or  mercury 
— is  present.  By  taking  account  of  the  differences  in  the 
properties  of  these  chlorides  it  is  not  difficult  to  decide  of 
which  of  them  a  precipitate  consists. 

HYDROXIDES. 

Experiment  203. — To  some  pieces  of  freshly-burnt  lime 
add  enough  cold  water  to  cover  it.  The  action  which  takes 
place  is  represented  by  the  equation 

CaO  +  H20  =  Ca(OH),. 

The  process  is  known  as  slaking. 

Experiment  204. — To  a  small  quantity  of  a  dilute  solution 
of  magnesium  sulphate  add  a  dilute  solution  of  caustic  soda. 
The  white  precipitate  is  magnesium  hydroxide.  [Would  you 

*  There  are  two  chlorides  of  mercury.  Only  one  of  them,  mercurous 
chloride,  is  insoluble. 


814    EXPEE1MENTS  TO  ACCOMPANY  CHAPTER  XXIV. 

expect  this  precipitate  to  be  soluble  in  sulphuric  acid  ?  in  hy- 
drochloric acid  ?  in  nitric  acid  ?]  The  answers  follow  from 
these  considerations  :  When  acids  act  upon  hydroxides,  salts 
are  formed  ;  magnesium  sulphate  is  soluble,  as  is  seen  by  the 
fact  that  we  started  with  a  solution  of  this  salt ;  the  only  inso- 
luble chlorides  are  those  of  silver,  lead,  and  mercury  ;  all 
nitrates  are  soluble. 

When  a  solution  of  an  iron  salt  is  treated  with  sodium  hy- 
droxide a  precipitate  of  iron  hydroxide  is  formed  : 

FeCls  +  3NaOH  =  Fe03H3  +  SNaOl. 

Experiment  205. — To  a  dilute  solution  of  that  chloride  of 
iron  which  is  known  as  ferric  chloride  add  caustic  soda. 
The  reddish  precipitate  which  is  formed  is  ferric  hydroxide. 
[From  the  general  statements  made  above,  would  you  expect 
this  precipitate  to  be  soluble  in  hydrochloric  acid?  in  nitric 
acid  ?  Try  each.  Is  it  soluble  in  sulphuric  acid  ?] 

Experiment  206. — Add  to  a  solution  of  an  aluminium  salt 
sodium  hydroxide.  After  a  precipitate  is  formed  continue  to 
add  the  sodium  hydroxide.  Perform  similar  experiments  with 
a  chromium  and  with  a  lead  salt.  Boil  each  of  the  solutions 
obtained.  Treat  a  solution  of  copper  sulphate  with  sodium 
hydroxide  in  the  cold.  Heat. 


SULPHATES. 

Experiment  207  — Make  a  dilute  solution  of  barium  chlo- 
ride, of  lead  nitrate,  of  strontium  nitrate.  To  a  small  quan- 
tity of  each  in  a  test-tube  add  a  little  sulphuric  acid.  [What 
remains  in  solution  ?]  Make  a  somewhat  concentrated  so- 
lution of  calcium  chloride.  To  this  add  sulphuric  acid. 
[What  is  in  solution  ?]  Add  more  water,  and  see  whether  the 
precipitate  will  dissolve.  The  formulas  of  the  salts  used  in 
the  experiments  are  barium  chloride,  BaCl2 ;  lead  nitrate, 
Pb(N03)2  ;  strontium  nitrate,  Sr(N03)2.  [Write  the  equations 
expressing  the  reactions.  ]  If  to  the  solutions  of  the  salts  any 
soluble  sulphate  is  added  instead  of  sulphuric  acid,  the  same 
insoluble  sulphates  will  be  formed.  The  sulphates  of  iron,  cop- 
per, sodium,  and  potassium  are  among  the  soluble  sulphates. 
Make  dilute  solutions  of  small  quantities  of  each  of  these,  and 
add  them  successively  to  the  solutions  of  barium  chloride, 


REDUCTION  OF  SULPHATES  TO  SULPHIDES.       815 

lead  nitrate,  and  strontium  nitrate.  The  formula  of  iron  sul- 
phate is  FeS04 ;  of  copper  sulphate,  CuS04 ;  of  sodium  sul- 
phate, Na2S04 ;  and  of  potassium  sulphate,  K2S04.  Write  the 
equations  representing  the  reactions  which  take  place  in  the 
above  experiments.  It  need  hardly  be  explained  that  the 
action  consists  in  an  exchange  of  places  on  the  part  of  the 
metals.  Thus,  when  the  soluble  salt  iron  sulphate,  FeS04, 
is  brought  together  with  the  soluble  salt  barium  chloride, 
BaCl2,  the  insoluble  salt  barium  sulphate,  BaS04,  and  the 
soluble  salt  iron  chloride,  FeCl2,  are  formed  : 

FeS04  +  BaCl2  =  FeCl2  +  BaS04. 

REDUCTION  OF  SULPHATES  TO  SULPHIDES. 

Experiment  208. — Mix  and  moisten  a  little  sodium  sulphate 
and  finely-powdered  charcoal.  Heat  the  mixture  for  some 
time  in  the  reducing  flame.  After  cooling  scrape  off  the  salt, 
dissolve  it  in  a  few  cubic  centimeters  of  water,  and  filter 
through  a  small  filter.  If  the  change  to  the  sulphide  has 
taken  place,  sodium  sulphide,  Na2S,  is  in  solution.  A  solu- 
tion of  a  sulphide  when  added  to  a  solution  containing  copper 
gives  a  black  precipitate  of  copper  sulphide.  Try  this;  also 
try  the  action  on  the  solution  of  the  salt  of  copper  of  some  of 
the  sulphate  from  which  the  sulphide  was  made. 

CARBONATES. 

Experiment  209. — The  formation  of  carbonates  by  the  ad- 
dition of  soluble  carbonates  to  solutions  of  salts  of  metals  whose 
carbonates  are  insoluble,  is  illustrated  by  the  following  experi- 
ments: Make  solutions  of  copper  sulphate,  iron  sulphate,  lead 
nitrate,  silver  nitrate,  calcium  chloride,  barium  chloride. 
Add  to  each  a  little  of  a  solution  of  a  soluble  carbonate,  as 
sodium  carbonate,  potassium  carbonate,  ammonium  carbon- 
ate. Note  the  result  in  each  case.  Filter  off  all  the  pre- 
cipitates and  prove  that  they  are  carbonates.  This  may  be 
done  by  treating  them  with  dilute  acids,  which  decompose 
them,  causing  an  evolution  of  carbon  dioxide,  which  can  be 
detected  by  passing  a  little  .of  it  into  lime-water.  In  some  of 
the  cases  mentioned  the  insoluble  salts  formed  are  basic  car- 
bonates, as,  for  example,  those  of  copper  and  magnesium.  The 
salts  of  silver,  calcium,  and  barium  are  the  normal  carbonates 
Ag.COa,  BaC03,  and  CaC03. 


816    EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXV. 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXV. 
POTASSIUM  SALTS. 

Experiment  210. — In  preparing  potassium  iodide  from 
iodine  and  potassium  hydroxide,  proceed  as  follows  :  To  30 
grams  iodine  use  15  grams  hydroxide.  Dissolve  the  latter  in 
100  cc.  water.  Add  half  this  solution  to  the  iodine  in  a  por- 
celain evaporating  dish.  Now  slowly  add  the  rest  of  the 
liquid  until  the  color  disappears.  Concentrate  the  liquid  to 
a  syrupy  consistence,  add  1  gram  finely-powdered  charcoal, 
mix,  and  evaporate  to  dryness.  The  residue  is  then  heated 
to  redness  in  an  iron  vessel.  After  cooling  extract  with  water. 

Experiment  211. — Potassium  iodide  can  also  be  prepared 
by  the  following  method  :  Bring  together  in  a  capsule  200 
grams  water,  10  grams  iron  filings,  and  40  grams  iodine  ;  mix, 
and  heat  gently.  When  the  solution  has  become  green,  de- 
cant, filter,  and  wash.  Now  heat  the  liquid  nearly  to  boiling, 
and  gradually  add  a  solution  of  35  grams  potassium  carbonate 
in  100  grams  water.  Filter,  wash,  and  evaporate. 

Experiment  212. — Dissolve  50  grams  potassium  carbonate 
in  500  to  600  cc.  water.  Heat  to  boiling  in  an  iron  or  a  silver 
vessel,  and  gradually  add  the  slaked  lime  obtained  from  25 
to  30  grams  of  good  quicklime.  During  the  operation  the 
mass  should  be  stirred  with  an  iron  spatula.  After  the  solu- 
tion is  cool,  draw  it  off  by  means  of  a  siphon  into  a  bottle. 
This  may  be  used  in  experiments  in  which  caustic  potash  is 
required. 

Experiment  213. — Mix  together  15  grams  potassium  nitrate 
and  2.5  grams  powdered  charcoal.  Set  fire  to  the  mass. 

Experiment  214. — Treat  a  quantity  of  wood  ashes  with 
water.  Filter,  and  examine  by  means  of  red  litmus-paper. 
Evaporate  to  dryness.  What  evidence  have  you  that  the 
residue  contains  potassium  carbonate  ? 

SODIUM  SALTS. 

Experiment  215. — Make  a  supersaturated  solution  of  sodi- 
um sulphate  by  heating  an  excess  of  the  salt  with  water  at  33°. 
Filter  the  solution  into  small  flasks  and  cork  them.  On  re- 
moving the  corks  and  agitating  the  vessels,  the  salt  will  sud- 
denly crystallize  out. 

Experiment  216. — Pass  carbon  dioxide  into  a  strong  solu- 
tion of  ammonia  (about  100  cc.)  until  it  is  no  longer  absorbed. 


CALCIUM  SALTS.— MAGNESIUM  AND  ITS  SALTS.    817 

A  solution  of  acid  ammonium  carbonate  is  thus  obtained. 
Add  this  to  a  concentrated  solution  of  sodium  chloride  as 
long  as  a  precipitate  is  formed.  Filter  off  the  precipitate, 
and  dry  it  by  spreading  it  upon  layers  of  filter-paper.  Heat 
some  of  the  salt  when  dry,  and  determine  whether  the  gas 
given  off  is  carbon  dioxide  or  not.  When  gas  is  no  longer 
given  off  by  heat,  let  the  tube  cool  and  examine  the  residue. 

Experiment  217. — Make  ammonium  sulphide  thus  :  Divide 
a  given  quantity  of  a  solution  of  ammonia  into  two  equal 
parts.  Saturate  one  half  by  passing  hydrogen  sulphide  through 
it,  and  then  add  the  other  half. 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXVI. 

CALCIUM  SALTS. 

Experiment  218. — Dissolve  10  to  20  grams  of  limestone  or 
marble  in  common  hydrochloric  acid.  Filter,  and  evaporate 
to  dryness.  Expose  a  few  pieces  of  the  residue  to  the  air. 

MAGNESIUM  A$TD  ITS  SALTS. 

."Experiment  219. — Make  anhydrous  magnesium  chloride 
thus  :  Dissolve  180  grams  magnesia  usta  in  ordinary  hydro- 
chloric acid  ;  shake  the  solution  with  an  excess  of  magnesia 
to  remove  iron  and  aluminium ;  filter ;  add  400  grams  am- 
monium chloride ;  evaporate  to  dryness,  keeping  the  mass 
constantly  stirred.  The  double  salt  thus  formed  must  be 
dried  until  a  small  specimen  put  in  a  test-tube  is  found  not 
to  give  off  water  when  heated.  The  dry  salt  is  then  ignited 
in  a  crucible  placed  in  a  furnace  until  ammonium  chloride  is 
no  longer  given  off,  when  the  molten  mass,  which  is  anhydrous 
magnesium  chloride,  is  poured  out  on  a  stone  and,  after  it  is 
broken  up,  it  is  put  in  a  dry  bottle  provided  with  a  good 
stopper. 

Experiment  220. — Mix  6  parts  anhydrous  magnesium 
chloride,  1  part  of  a  mixture  of  sodium  and  potassium  chlo- 
rides, prepared  by  melting  the  two  together  and  breaking  up 
after  cooling,  1  part  powdered  fluor-spar,  and  1  part  sodium. 
Throw  this  mixture  all  at  once  into  a  red-hot  crucible  in  a 
furnace,  and  cover  the  crucible.  In  a  few  moments  a  curious 
sound  is  heard,  and  this  indicates  that  the  reaction  is  taking 
place.  Now  take  the  crucible  out  of  the  furnace,  and  stir  the 
liquid  in  it  with  the  aid  of  a  clay  pipe-stem.  This  causes  the 


818    EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXVIII. 

particles  of  the  metal  to  collect  in  one  large  spherical  mass. 
After  cooling,  break  the  crucible,  separate  the  metallic  ball 
from  the  slag,  and  wash  it  quickly  with  hydrochloric  acid  to 
remove  superficial  impurities.  If  the  slag  is  melted  with  a 
quarter  the  weight  of  sodium  that  was  used  at  first,  a  second 
smaller  piece  of  magnesium  will  be  obtained. 


EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXVII. 
ALUMINIUM  CHLORIDE. 

Experiment  221. — Aluminium  chloride  is  made  thus : 
Mix  aluminic  oxide  with  starch-paste  ;  form  the  mass  into 
small  balls  of  the  size  of  ordinary  marbles  ;  ignite  these  in  a 
crucible  in  a  furnace  ;  put  them  in  a  porcelain  tube,  and  then 
pass  dry  chlorine  over  them,  at  the  same  time  heating  the  tube 
to  redness.  The  chloride  will  sublime  in  the  front  end  of  the 
tube  or  in  a  receiver  if  the  heat  is  sufficient.  It  can  be  puri- 
fied by  subliming  it  over  heated  iron  or  aluminium. 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXVIII. 
COPPER  AND  ITS  SALTS. 

Experiment  222. — Cuprous  chloride  is  best  made  as  fol- 
lows :  Saturate  a  solution  of  1  part  sodium  chloride  and  2| 
parts  crystallized  copper  sulphate  with  sulphur  dioxide.  Fil- 
ter, and  wash  with  acetic  acid. 

SILVER  AND  ITS  SALTS. 

Experiment  223. — Dissolve  a  ten  or  twenty-five  cent  piece 
in  dilute  nitric  acid.  Dilute  the  solution  to  200  to  300  cc. 
with  hot  water.  Add  a  hot  solution  of  common  salt  until  it 
ceases  to  produce  a  precipitate.  Filter  off  the  white  silver 
chloride  and  wash  with  hot  water.  Dry  the  precipitate  on 
the  filter,  by  placing  the  funnel  with  the  filter  and  precipitate 
in  an  air-bath  heated  to  about  110°.  Eemove  the  precipitate 
from  the  filter  and  put  it  into  a  porcelain  crucible.  Heat 
gently  with  a  small  flame  until  the  chloride  is  melted ;  then 
let  it  cool.  Cut  out  a  piece  of  sheet  zinc  large  enough  to 
cover  the  bottom  of  the  crucible,  and  lay  it  on  the  silver 
chloride.  Now  add  a  little  water  and  a  few  drops  of  dilute 


ZINC  AND  ITS  SALTS— TIN  AND  ITS  COMPOUNDS.   819 

sulphuric  acid,  and  let  the  whole  stand  for  twenty-four  hours. 
The  silver  chloride  is  reduced  to  silver,  and  zinc  chloride  is 
formed : 

Zn  +  SAgCl  =  ZnCl2  +  2Ag. 

Take  out  the  piece  of  zinc  and  wash  the  silver  with  a  little 
dilute  sulphuric  acid,  and  then  with  water.  Dissolve  the 
silver  in  dilute  nitric  acid  and  evaporate  to  dryness  on  the 
water-bath,  so  that  all  the  nitric  acid  is  driven  off.  Dissolve 
the  residue  in  water,  and  put  the  solution  either  in  a  bottle 
of  dark  glass  or  one  wrapped  in  dark  paper. 

Experiment  224. — To  a  solution  of  silver  nitrate  contain- 
ing about  5  grams  of  the  salt  in  100  cc.  water,  add  a  few  drops 
of  mercury,  and  let  it  stand.  In  a  few  days  the  silver  will  be 
deposited  in  the  form  of  delicate  crystals.  The  formation  is 
called  the  "silver  tree." 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXIX. 

ZINC  AND  ITS  SALTS. 

Experiment  225. — Heat  a  small  piece  of  zinc  on  charcoal 
in  the  oxidizing  flame  of  the  blow-pipe.  The  white  fumes  of 
zinc  oxide  (philosopher's  wool)  will  be  seen,  and  the  charcoal 
will  be  covered  with  a  film  which  is  yellow  while  hot,  but  be- 
comes white  on  cooling. 

Experiment  226. — Dissolve  some  zinc  dust  in  a  solution  of 
sodium  hydroxide,  and  see  whether  hydrogen  is  given  off. 

MEKCURY  AND  ITS  SALTS. 

Experiment  227. — Make  a  solution  of  mercurous  nitrate 
by  treating  at  the  ordinary  temperature  an  excess  of  mercury 
with  nitric  acid,  which  is  not  too  concentrated  ;  and  with  this 
solution  study  the  conduct  of  mercurous  salts. 

Experiment  228. — Heat  some  of  the  solution  of  mercurous 
nitrate  to  boiling,  then  add  a  few  drops  of  concentrated  nitric 
acid,  and  boil  again.  With  the  solution  thus  obtained  study 
the  conduct  of  mercuric  salts. 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXX. 
TIN  AND  ITS  COMPOUNDS. 

Experiment  229. — Dissolve  tin  in  hydrochloric  acid  and  let 
the  product  crystallize. 


820    EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXXI. 

Experiment  230. — Pass  dry  chlorine  over  granulated  tin 
contained  in  a  retort  connected  with  a  receiver,  using  the  ar- 
rangement illustrated  in  Fig.  64,  page  800.  Kedistil  the  prod- 
uct. Treat  some  of  the  liquid  with  water,  and  boil. 

LEAD  A^D  ITS  COMPOUNDS. 

Experiment  231. — Make  specimens  of  lead  chloride  and 
lead  iodide,  and  crystallize  them  from  water. 

Experiment  232. — Make  lead  sesquioxide  by  bringing  to- 
gether lead  acetate  and  sodium  hydroxide,  and  treating  the 
solution  with  a  solution  of  sodium  hypochlorite. 

Experiment  233. — Treat  some  red  lead  with  dilute  nitric 
?icid.  Filter,  wash,  and  treat  the  substance  left  on  the  filter 
with  hydrochloric  acid. 


EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXXI. 
CHROMIC  ACID  AND  THE  CHROMATES. 

Experiment  234. — Powder  some  chromic  iron  very  finely. 
Add  3  grams  to  a  molten  mixture  of  3  grams  each  of  potas- 
sium carbonate,  potassium  hydroxide,  and  potassium  nitrate, 
heated  in  a  porcelain  crucible.  After  cooling  treat  the  mass 
with  water.  Potassium  chromate  is  in  the  solution. 

Experiment  235. — To  the  solution  of  potassium  chromate 
obtained  in  the  last  experiment  add  nitric  acid  to  decompose 
the  unacted-upon  potassium  carbonate,  and  give  the  solution 
an  acid  reaction.  The  color  will  change  from  yellow  to  red. 
The  red  color  indicates  the  presence  of  the  dichromate. 

Experiment  236. — Treat  a  solution  of  10  to  20  grams 
potassium  dichromate  with  potassium  hydroxide  ,  until  the 
color  becomes  pure  yellow,  and  evaporate  to  crystallization. 

Experiment  237. — Make  a  solution  of  potassium  dichro- 
mate saturated  at  the  ordinary  temperature.  Pour  into  this 
1£  times  its  volume  of  ordinary  concentrated  sulphuric  acid. 
After  the  liquid  cools,  and  the  chromium  trioxide  separates, 
filter  with  the  aid  of  a  filter-pump  through  glass-wool. 

Experiment  238. — To  a  solution  of  potassium  dichromate 
add  some  hydrochloric  acid  and  a  little  alcohol.  On  boiling, 
the  alcohol  is  oxidized,  and  the  solution  now  contains  chromic 
chloride. 


MANGANESE  AND  ITS  COMPOUNDS— PLATINUM.     821 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXXII. 

MANGANESE  AND  ITS  COMPOUNDS. 

Experiment  239. — Make  and  crystallize  some  manganous 
chloride  by  treating  manganese  dioxide  with  hydrochloric 
acid.  Also  make  some  manganous  sulphate  by  heating  man- 
ganese dioxide  with  sulphuric  acid.  Use  these  solutions  for 
the  purpose  of  studying  the  conduct  of  manganous  salts. 

Experiment  240. — In  a  small  porcelain  crucible  heat  to- 
gether 5  grams  manganese  dioxide,  5  grams  solid  potassium 
hydroxide,  and  2J  grams  potassium  chlorate.  When  the 
mass  has  turned  green,  dissolve  the  contents  in  water  and 
neutralize  most  of  the  free  alkali  in  the  solution.  Or  pass 
carbon  dioxide  through  the  solution  without  boiling. 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXXIII. 
IRON  AND  ITS  COMPOUNDS. 

Experiment  241. — Make  ferric  chloride  by  heating  the 
purest  iron  wire  in  a  current  of  chlorine.  Also  make  a  solu- 
tion by  dissolving  iron  in  hydrochloric  acid,  and  oxidizing  the 
solution  with  nitric  acid. 

Experiment  242. — Make  ferrous  sulphate  by  dissolving  iron 
in  dilute  sulphuric  acid,  and  evaporating  to  crystallization. 
Dissolve  equivalent  quantities  of  ferrous  sulphate  and  ammo- 
nium sulphate,  and  evaporate  to  crystallization. 

Experiment  243. — Dissolve  ferric  hydroxide  in  sulphuric 
acid,  and  evaporate  to  dryness. 

EXPERIMENTS  TO  ACCOMPANY  CHAPTER  XXXIV. 
PLATINUM. 

Experiment  244. — Prepare  a  solution  of  platinic  chloride 
as  follows  :  Heat  platinum  in  a  flask  with  concentrated  nitric  • 
acid,  adding  from  time  to  time  a  few  drops  of  hydrochloric 
acid.     After  the  metal  is  dissolved  evaporate  to  dryness.    Dis- 
solve in  water  and  filter. 

CONCLUSION. 

At  the  end  of  each  chapter  treating  of  the  metallic  or  base- 
forming  elements  there  is  given  a  list  of  such  reactions  as  are 


822  INORGANIC  CHEMISTRY, 

of  special  value  for  analytical  purposes,  together  with  such  ex- 
planatory statements  as  seem  called  for.  The  student  who  is 
engaged  in  analytical  work  will  generally  find  these  explana- 
tions sufficient  to  enable  him  to  keep  his  ideas  clear  in  regard 
to  the  reactions  with  which  he  is  dealing,  provided,  at  the 
same  time,  he  carefully  studies  the  chapter  to  which  the  ex- 
planations form  an  appendix.  As  an  introduction  to  analyti- 
cal work,  a  general  study  of  chemical  reactions  is  necessary, 
and  the  fuller  this  is  the  better.  There  are  many  small  books 
in  existence  in  which  good  directions  are  given  for  work  of 
this  kind.  The  student  is,  however,  advised  to  supplement 
the  book  he  may  be  using  by  such  experiments  as  may  suggest 
themselves  on  reading  the  corresponding  chapters  in  this 
book.  The  directions  there  found  will  generally  be  quite  suf- 
ficient for  the  purpose,  and  it  is  therefore  not  considered 
necessary  to  give  more  specific  directions  in  this  place.  In 
general,  the  more  the  student  occupies  himself  in  the  labora- 
tory with  chemical  substances  the  more  rapidly  will  his  chemi- 
cal ideas  grow.  But  it  is  necessary  that  he  should  avoid 
working  by  "rule  of  thumb,"  and  for  this  purpose  constant 
reference  to  some  larger  text-book  in  which  the  relations  be- 
tween the  substances  and  the  reactions  he  is  dealing  with  are 
discussed  in  a  broad  way  is  of  the  highest  importance. 


APPENDIX  II. 

THE  following  tables  have  been  prepared  to  furnish  those 
who  may  use  this  book  as  a  reference-book  in  connection  with 
laboratory  work  with  such  information  as  they  may  need,  but 
which  they  can  perhaps  get  only  from  works  not  always  easily 
available.  It  was  not  considered  advisable  to  include  tables 
required  in  gas  analysis  and  general  quantitative  analysis. 
For  part  of  the  material  I  am  indebted  to  that  excellent 
work,  the  Chemiker-Kalender,  by  E.  Biedermann.  In  the 
tables  of  solubilities  the  substances  have  been  arranged 
alphabetically  as  salts  of  acids,  as  for  example  the  acetates, 
chlorides,  sulphates,  and  sulphydrates. 

J.  ELLIOTT  GILPIK. 


824 


APPENDIX  II. 


ATOMIC  WEIGHTS. 
(From  "  The  Constants  of  Nature,"  by  F.  W.  Clarke,  1897.) 


H  =  l 
2691 

O  =  16 
27.11 

Molybdenum.  .  .  . 

H  =  1 
95.26 

Antimony    •  •  • 

119  52 

12043 

Neodymium  

139.70 

9 

? 

Nickel  

58  24 

74  44 

75  01 

Nitrogen 

13  93 

Barium    .  .  .  .... 

136  39 

137  43 

189  55 

Bismuth 

206  54 

208  11 

Oxygen 

15  88 

10  86 

10  95 

Palladium        .  .  . 

105  56 

Bromine   •  •  •       . 

79  34 

79  95 

Phosphorus    .... 

3079 

111  10 

111.95 

Platinum  

193  41 

131  89 

132.89 

Potassium  

38.82 

Calcium.       .  .... 

39  76 

40  07 

Praseodymium    . 

142  50 

Cirbon        

11  92 

12  01 

Rhodium  

102  23 

Cerium          

139  10 

140.20 

Rubidium  

84  78 

Chlorine 

35  18 

35  45 

Ruthenium  .  .  . 

100  91 

Chromium 

51  74 

52  14 

Samarium  

149  13 

Cobalt       .            . 

58  49 

58  93 

43  78 

Columbium 

93  02 

93  73 

Selenium         .  .  .  . 

78  42 

CoDoer  . 

63.12 

63.60 

Silicon  

28.18 

Erbium  .    . 

165  06 

166  32 

Silver  

107.11 

18.91 

19.06 

22.88 

Gadolinium 

155  57 

156  76 

Strontium     . 

86  95 

Gallium 

69  38 

69  91 

Sulphur  

31  83 

Germanium  

71  93 

72.48 

Tantalum  

181  45 

9.01 

9  08 

126  52 

Gold 

195  74 

197  23 

Terbium  

158  80 

? 

? 

Thallium  

202.61 

Hydrogen 

1  000 

1  008 

Thorium 

230  87 

Indium 

112  99 

113  85 

Thulium  

169  40 

Iodine  ... 

125  89 

126.85 

Tin         

118  15 

Iridium 

191  66 

193  12 

Titanium          . 

47  79 

Iron 

55  60 

56.02 

Tungsten 

183  43 

Lanthanum 

137  59 

138.  '14 

Uranium 

237  77 

Lead  

205.36 

206.92 

50  99 

Lithium 

6  97 

7  03 

Ytterbium 

171  88 

Magnesium  

24  10 

24.28 

Yttrium     

88  35 

Manganese 

54  57 

54  99 

Zinc 

64  91 

Mercury.. 

198.49 

200.00 

Zirconium  .  . 

89.72 

O  =  16 

95.99 

140.80 

58.69 

14.04 

190.99 

16.00 

106.36 

31.02 

194.89 

39.11 

143.60 

103.01 

85.43 

101.68 

15026 

44.12 

79.02 

28.40 

107.92 

23.05 

87.61 

32.07 

182.84 

127.49 

160.00 

204.15 

232.63 

170.70 

119.05 

48.15 

184.83 

239.59 

51.38 

173.19 

89.02 

65.41 

90.40 


APPENDIX  II. 


825 


MELTING-POINTS  AND   BOILING  POINTS  OF  THE 

ELEMENTS. 

Where  the  values  of  different  observers  vary,  the  highest  and  lowest 
ill  be  given. 


Melting-point. 

Boiling-point. 

Aluminium.  ..... 
Antimony  

600-850 
425-450 

n.v.  at  white  heat 
1090-1700 

—  189  6 

—  186  9 

Al'senic              .    . 

under  pressure  at  red  heat 

sublimes  at  450 

Barium  

red  heat 

Beryllium 

below  1000 

Bismuth.  .  .       ... 

260-270 

1090-1450 

Boron       

Bromine  

—  7.3 

59-63 

Cadmium 

310-320 

760-860 

CflBsiuni          ...» 

26  5 

Calcium  

red  heat 

n  v 

infusible 

Cerium 

between  Sb  and  Ag 

Chlorine  

solidifies  at  —  102 

—  33.6 

Chromium  

higher  than  platinum. 

1500-1800 

Copper 

1000-1330 

Fluorine  

Gallium  * 

30.15 

Germanium  

about  900 

Gold         

1035-1250 

Indium  

176 

red  heat 

Iodine  

113-115 

above  200 

Indium 

1950-2500 

Iron    pure.  . 

1500-1800 

322-335 

between  1450-1600 

Lithium 

180 

Magnesium    
Manganese  

500-800 
1600 

about  1100 

Jdercury 

—  38  50  to  —  40  5 

357  25 

Molybdenum  
Nickel  

does  not  melt  at  a  white  heat 
1450-1600 

Nitrogen  ., 

solidifies  at  —  203  at  60-70  mm. 

—  1944 

Osmium     .     «... 

2500 

Oxveen  . 

below  —  21]  5  at  9  mm 

—  181  4 

Palladium  

1360-1950 

Phosphorus     . 

44  2 

287-290 

Platinum  

1460-2200 

Potassium 

58  62  5 

667  731 

Rhodium.  .    . 

2000 

Rubidium  

38  5 

Selenium           .  .  . 

217 

bet  676  and  683 

Silicon  .... 

about  1400 

Silver  

954-1040 

Sodium 

90  97  6 

742-954 

Strontium.  . 

red  heat 

Sulphur,  rhombic, 
mono.  .. 
'  '        amor.  .  .  . 
Tellurium  
Tin  

113-115 
120 
above  120 
425-525 
226  5-235 

447-448.4 
bet  1450  and  1600 

Zinc  

412-420 

891  1040 

826 


APPENDIX  II. 


ttj 

53*3 


-:-:dH  •  -£o 


dgS 


=>.  sw 


^  I 


Melting'- 


OOi  W 
<M  CO  ~ 


o   1 


bfri 

o 


o    o, 


»  O 
?    cS 


W    B        55 


te 

e 


at 
ce 


. 


..fi|  Iss^l 

ll£'"V 


o    .iJ 


cid. 
acid 


a  ^  ^  ::  o 


Sodium 
Urauyl 


2    1§SJ 

1  Sill 
ilgfS^ 


APPENDIX  II. 


827 


• 


I   5 


§    - 


t-sfii. 

§§S§ 


88888 


8 


«;    -5 


888880'    'S    |    ^ 


I* 


J 

O*T3 

T-t 

a 


i§ri  § 
:!«  I 

I     QJ  ,Q 


828 


APPENDIX  II. 


flj 

*$t 

'  *  *"  d 

»• 


80^ 


o  » 

(D  - 

a 


Sao  -     . 
is^-*!1* 


•o'B     1o 


o    .3  a 


O     ®  in 

8  o^ 

^»  ^^-r-J 


£10   £ 


•  t  s   I 

,      O  TH  O 

!     O  \  O 


00  0°° 

00  05  fc.'-JrH 

JH  ^  a>0'-'co 

C^  ^^  *^          *H  CO 

^^  ^D  03                    • 

.  .  >           i-H  T-H 

O  O  P              C5 


5    8  SS5 


I? 

•5 

•3 

S 


Oq    o 

H?W         W 


*    -0 
O    PH 


B8 


O  " 


iS.'*5  ?"  a  ® 
sill S? 


s  S-s  ^ 

si» «  "-'S  S  ^   u. 

O   S3  OQ  S  ^ 

S-IS3  §: 

M    C2  rt 


O    <U  .S         .S  t- 

7='C    -          t-          O 

"^•r  o 


c  c^s     ^ 


3  a 

o  b^ 53    6 

gWP,g«*1   <1       <j         «X^ 


§         111 

a  s   "=  o  § 
•s      ii  I 


APPENDIX  II. 


829 


cohol,  Acids, 
lkalies,  and 
Other  Sol- 
vents. 


s  § 


ft 


o    S 

03        O 
•§ 


i 


-a  i  i* 


£? 


\t 

r~  £ 


i  .Si  :£ 


W 


SE  S 

13    '   13 


q  q 

as 


«  e« 

0  Q 

T3  03 

O  O 


J  Q 
Q 


o    o 


00 

ww 


-  B      .8J., 


o    fe 


?l 
II 
II 

il 


a 


e 
chlori 
chlor 


"  o>  fl> 

O>    Q^  I  T3  rO 

CJ  ^  T3       0>  -  "^ 

?S"S  ? 

3-31  a  !p?»s 

o  »  „  -3  gg^  ggg 
.Sig  §  ^3^3-sS 
BBS  o  «-=>«« 

III    I 

060  o 


o 
^ 
'C 

o 

1 

• 
3 
c 


osphorus 


num  chloride 
ssium  (Chlorid 


II 

B-3 


£T£ 


GQ 


830 


APPENDIX  II. 


& 


•   • 


ww 


2    2 

8O 
O 
C2       O 


a  :^ 


s      ~ 


§ 


M^l 


G 
O 

I? 

•-3 


§     § 


§ 


•3        ° 

»S    w 

7"  i 


O 

s 

a  §  §5iN 

OQ       .S        "<*&¥>     ^ 


O        j 


^    o 

S   « 

^         eo 


o    o 

OQ       CO 


O 

-r^p?   o9' 


HLORIDES  (cont'd)  : 
Silicon  chloroform 


.    .    . 

.        .        . 

S 

s 

c3 

•  *r^          . 

2 

O 

CJS       4 

S 

j 

!       C 

« 

1 

»£»% 
?-§Sb 

Is 

O    rf    - 

<uS^  ^ 

•c  5-  •?  'S 

"C^^  o 

0,0  j  3 

•^  o  o  7s 


o    .-  o 


o     ^    S-^ 


um 
um 


l 

h 
e 


^3      0^3 

l',t 

H 

c^   eg 


ori 
ac 
ri 
lor 


l 
x 
lo 
h 
e 


m  c 
n  h 
c 


Tha 
Tun 
Uram  h 
Vanadium  c 
Zinc  chlorid 


TE 


-°  S^ 

s  2 1 

BD-3-gP 

1a« 
Sl-SiS 

gl^^ 

Q 


APPENDIX  II. 


831 


a  co  ~  ^—'  ,— "       d  a ' 


13  1 


§  ;  _ 

'p.l.S 
£ 


§ 


!, 

.s° 


O        »O  O  »O  GO 
T-I        00-^QOTH 

I  I      I      I      I 


. 

F 


s 


>> 


0 

s  § 
ft 


<S         o' 

^      O 


®        r5 


tf 

"Ili 

00 


5  A  5 


832 


APPENDIX  II. 


So  :    s 

1  O     . 


«   *  rj 

:  8     a  S 


oo    ;      o 
01    •      g 


.i 


a 
a. 


8    . 

T-l          CU 


W 


1 


I    a 

o      o 

g    g 


:H  > 


}>i— i  (_HH    {>• 


I ooooo 

«  ^-^   C^   X ^>— '^7 


o- 

HH  P 


cont'd} 
ydrox 


i  i  :  :S 

4i-§  i§ 

gill' 

'H§a 

Ms-ll 


roxide 

ydroxide. 
ydroxide. 
roxide 
ydroxide. 


de 

S  : 

chlorite 
pochlori 


! «  t  ; 

.  [O 

'  T3  oJ     •' 

03 .2 12   * 

^5  S  o  oi 

O   3'~T3 

11  §1 
11  tig 

"    CJ  _rt    o  •— c 

0^£(£X 


: 

mangan 


ANGANAT 

Potassiu 


APPENDIX  II. 


833 


fl. 


F 


~o      .  .0  -    . 

' 


g    o 


•so 


iti  I* 


^        b 


\a 

<M   O 


A'      ^       a        S  *• 

O          g,  .„  ^, 


H 

'!. 


a 

a. 
UP 

C  o 

s 
"S 


a 


g    co    ^ 


w 


(53 


ooo 

C*       CT. 


+   o 


+1 !  8 


PQQO 


Q>        '      ' 

l    iig 


^ 


'P^§sg§ 

^  i-T!  S  C  ^  H 


a  I 

w -5 

^<j  -^pppqpq 

O 


.2  g  a 

s  3 


2  d 

2  I 

83  Oj 

o  o 


834 


*g. 

3«S 
<-r« 


Al 
O 


APPENDIX  II. 

*       M&rl  ^O 


* 


ft    . 


-  .2 


I* 


g  o£  8 

-H     <M     ,      S 


OHCC 


S  oj 
o-S 


S  w'6  (n-S-c  s  g^^^S  o  SS          ^^    g    as 

Hl|llftllli|S  i  &-s  s  Pi 

r,SS><0ttrS^^a*ooa    'x     3®    :,     5 'S 


wcocJ^rf-3  hfiD      3Q.S 

si  1111 1 II  H| 

!^{5  o  P^S 


APPENDIX  II. 


835 


acu    no 

a  a    " 

99     oo* 


:  a   -s 


og.2 


a  o 

o    ^** 

T3 


|d 
GO 

13 


d^    °°- 

a  i  - 


A       G.       O. 

a    a    a 


•3 
•s^i 

i*i 


i  1 


836 


APPENDIX  II. 


...§ 


ta 


3  i 

»      H 


s  °'  II 

°    *    7>-o 
"   .2   %<*. 


c 

fe 


I' 


c 

a. 
I? 

I 


*8 


>&  & 


-*2  "     oD 

«  •<  o 

^3  Q^3 

PH  PQ    P, 


.2  a 


1 


H 


§ 


., 

sill 


§    PL, 


ia 

-  £ 


,5  a 


I* 
It 


i 


OD       OQ 


APPENDIX  II. 


837 


fc 


IN 


.3        -3 


1? 


:oq  oq 

J3S  £5 


rn    -4-j  ~*j?  gg  ^ 

«  oqqo 
ggSS 


3    its 
o    :o 

5  jg" 


-  s  t 


••! 

?•§,  5 
«'3    S 


. 
§1 


I   lit  I    1 

rf        Ct    t>^rp3      *~^        rl 


O 


1« 

tl 

-2  >> 

^•^5 

."3  ^ 


§•83 


OQ 


<«JPH 


838 


APPENDIX  II. 


Alcohol,  Acids, 
Alkalies,  and 
Other  Sol- 
vents. 


«•  S  8  i 

o    .      a 


i® 


§£      7^      g 


c 

0 

* 


'o 

go 


O  !<«, 

T-l        TH  <M 

q  O  q 

W  £  W 


1 


+ 


1 

H« 

S3 

NH       ffi 

MJ         r+4 

«| 

CQ 

M 


f  sp-sil  f 

•  •  3       2  ^5  rH  Jd  ~^     M 

•  «  8*SB-M'a.l9i  fi 


ulph 
ide 
e 


ocy 
yau 

l  t 


5s 

£^3 


ill 


2  *  d 


QQ 


" 


33     v>2<2% 

afli-si5lsa 

2  tl5iill1 

^.«.2^o'SS2-S^So 

PnCCtSJ  gPn^  g<lPn  §&4^ 


otass 
ver 
c  s 


QQ 


02 


"^  *•    •  P-i  a 

"all 

S  o'S 
E^ 

H 


APPENDIX  II. 


839 


WEIGHT  OF  1   LITER  OF  GAS  AT  0°   AND  760  MM. 
PRESSURE   IN  LATITUDE  45°. 


Name. 

Formulae. 

Molecular  Weight. 

Weight  of  1  Liter, 
in  Grains. 

CH 

2584 

1  1611 

Ammonia       •     .  .  .      *    • 

NH3 

16  93 

0  7607 

Carbon  dioxide  

CO2 

43  68 

1  962 

"       monoxide  

CO 

Clo 

27.80 
70  36 

1.249 
3  1617 

Hydrochloric  acid.  .  .  . 

HC1 

36  18 

1  625 

H2 

2 

0  089873 

CH4 

15  92 

0.7153 

Nitric  oxide  

NO 

29  81 

1  339 

Nitrogen  

No 

2786 

1  252 

NO2 

45  69 

2053 

Nitrous  oxide 

N2O 

43  74 

1  9655 

Oxygen 

Oo 

31  76 

1  4^9 

Sulphur  dioxide  

SO2 

63  59 

28575 

Morley's  values  for  oxygen  and  hydrogen  were  adopted,  and  the  others 
calculated  on  this  basis. 


APPROXIMATE  COMPOSITION  OF  A  NUMBER  OF  ALLOYS. 


Pro] 

jortion 

sin 

ivhic 

h  Con 

ibim 

ition  t 

akes 

Place. 

Cu 

Zn 

Sn 

Ni 

Ag 

P 

Cd 

Bi 

Pb 

Al 

Sb 

Brflss        ......... 

66 

33 

3 

Gun-metal  

90 

10 

Bell  metal  
German  silver  
Silver  coins  .  .  . 

77 
60 
10 

20 

23 

20 

90 

Wood's  metal  
Rose's  metal  

13.5 
25 

10 

49.8 
50 

26.7 
25 

Britannia  metal  
Pewter  

7-8 
1 

2 

77-90 
2-20 

10-15 

Aluminium  bronze 
Art  bronze     .  . 

90 
86  6 

3  3 

6.6 

3  3 

10 

Soft  solder  

50 

50 

Babbitt  metal  
Nickel  coins  .  .    .. 

3.7 

75 

88.9 

95 

7.4 

Phosphor  bronze  .. 

90 

9 

0.5 

840  APPENDIX  II. 


FREEZING-MIXTURES. 

The  mixtures  given  below  will  be  found  useful  in  cases  where  low 
temperatures  are  desired. 

Temperature  falls 
from  to 

1  pt.  potassium  sulphocyanate  -f-  1  pt.  water -|-  18          —  21 

1  pt.  sodium  chloride  +  3  pts.  snow —  21 

3  pts.  crystallized  calcium  chloride  -j-  1  pt.  snow —  48.5 

1  pt.  snow  -f-  1  pt.  dilute  sulphuric  acid. ......    -f  5            —  41 

1  pt.  potassium  chloride  +  4  pts.  water. ./ —  11.8 

3  pts.  sodium  nitrate  +  4  pts.  water -(-  13.2       —  5.3 

Solid  carbonic  acid  +  ether —  100 


TABLE  OF  WEIGHTS  AND  MEASURES. 

WEIGHT. 

1  gram  =  15.432  grains ; 

1  grain  =    0.06479  grams ; 

1  ounce  avoirdupois  =  28.3495  grams  ; 

1  ounce  troy  =  31.1035  grams  ; 

1  pound  avoirdupois  =    0.45359  kilogram  ; 

1  pound  troy  =    0.37324  kilogram  ; 

1  kilogram  =    2.2046  pounds  avoirdupois. 

LENGTH. 

1  millimeter  =        0.03937  inches ; 
linch  =      25.3995  millimeters; 

1  meter          —        3.280899  feet ; 
1  foot  =        0.30479  meter ; 

1  mile  =  1609.31  meters. 

VOLUME. 

1  cubic  centimeter  —    0.061027  cubic  inches  ; 
1  cubic  inch  =  16  38,6  cubic  centimeters  ; 

1  cubic  centimeter  =    0.03519  fluid  ounce  ; 
1  fluid  ounce  =  28.41  cubic  centimeters ; 

1  liter  =    1.76  pints  ; 

1  liter  =    0.2201  gallon; 

1  gallon  =    4.543  liters. 


APPENDIX  II. 


841 


COMPARISON  OF  THE  TWADDELL  SCALE  WITH  THE 
BAUMft  AND   GAY-LUSSAC  SCALES. 


Tw. 

B. 

G.-L. 

Tw. 

B. 

G.-L. 

Tw. 

B. 

G.-L. 

Tw. 

B. 

G.-L. 

0 

0 

1.000 

44 

26.0 

1.220 

88 

44.1 

1.440 

131 

57.1 

1.655 

1 

0.7 

1.005 

45 

26.4 

1.225 

89 

444 

1.445 

132 

57.4 

1.660 

2 

1.4 

1.010 

46 

26.9 

1.230 

90 

44.8 

1.450 

133 

57.7 

1.665 

3 

2.1 

1.015 

47 

27.4 

1.235 

91 

45.1 

1.455 

134 

57.9 

1.670 

4 

2.7 

1.020 

48 

27.9 

1.240 

92 

45.4 

1.460 

135 

58.2 

1.675 

5 

3.4 

1.025 

49 

28.4 

1.245 

93 

45.8 

1.465 

136 

584 

1.680 

6 

4.1 

1.030 

50 

28.8 

1.250 

94 

46.1 

1.470 

137 

58.7 

1.685 

7 

4.7 

1.035 

51 

29.3 

1.255 

95 

46.4 

1.475 

138 

58.9 

1.690 

8 

5.4 

1.040 

52 

29.7 

1.260 

96 

46.8 

1.480 

139 

59.2 

1.695 

9 

6.0 

1.045 

53 

30.2 

1.265 

97 

47.1 

1.485 

140 

59.5 

1.700 

10 

6.7 

1.050 

54 

30.6 

1.270 

98 

47.4 

1.490 

141 

59.7 

1.705 

11 

7.4 

1.055 

55 

31.1 

1.275 

99 

47.8 

1.495 

142 

60.0 

1.710 

12 

8.0 

1.060 

56 

31.5 

1.280 

100 

48.1 

1.500 

143 

60.2 

1.715 

13 

8.7 

1.065 

57 

32.0 

1.285 

101 

48.4 

1.505 

144 

60.4 

1  720 

14 

9.4 

1.070  | 

58 

32.4 

1.290 

102 

48.7 

1.510 

145 

60.6 

1.725 

15 

10.0 

1.075 

59 

32.8 

1.295 

103 

49.0 

1.515 

146 

60.9 

1.730 

16 

10.6 

1.080  I 

60 

33.3 

1.300 

104 

49.4 

1.520 

147 

61.1 

1.735 

17 

11.2 

1.085  i 

61 

33.7 

1.305 

105 

49.7 

1.525 

148 

61.4 

1.740 

18 

11.9 

1.090 

62 

34.2 

.310 

106 

50.0 

1.530 

149 

61.6 

1.745 

19 

12.4 

1.095 

63 

34.6 

.315 

107 

50.3 

1.535 

150 

61.8 

1  750 

20 

13.0 

1.100 

64 

35.0 

.320 

108 

50.6 

1.540 

151 

62.1 

1.755 

21 

13.6 

1.105  ! 

65 

35.4 

.325 

109 

50.9 

1.545 

152 

62.3 

1.760 

23 

14.2 

1.110  ! 

66 

35.8 

.330 

110 

51.2 

1.550 

153 

62.5 

1.765 

23 

14.9 

1.115 

67 

36.2 

.335 

111 

51.5 

1.555 

154 

62.8 

1.770 

24 

15.4 

1.120 

68 

36.6 

.340 

112 

51.8 

1.560 

155 

63.0 

1.775 

25 

16.0 

1.125 

69 

37.0 

.345 

113 

52.1 

1.565 

156 

63.2 

1.780 

26 

16.5 

1.130 

70 

37.4 

1.350 

114 

52.4 

1.570 

157 

63.5 

1.785 

•27 

17.1 

1.135 

71 

37.8 

1.355 

115 

52.7 

1.575 

158 

63.7 

1.790 

28 

17.7 

1.140 

72 

38.2 

1.360 

116 

53.0 

1.580 

159 

64.0 

1.795 

29 

18.3 

1.145 

73 

38.6 

1.365 

117 

53.3 

1.585 

160 

64.2 

1.800 

30 

18.8 

1.150 

74 

39.0 

1.370 

118 

53.6 

1.590 

161 

64.4 

1.805 

31 

19.3 

1.155  ! 

75 

39.4 

1.375 

119 

53.9 

1.595 

162 

64.6 

1.810 

32 

19.8 

1.160 

76 

39.8 

1.380 

120 

54.1 

1.600 

163 

64.8 

1.815 

33 

20.3 

1.165 

77 

40.1 

1.385 

121 

54.4 

1.605 

164 

65.0 

1.820 

34 

20.9 

1.170 

78 

40.5 

1.390 

122 

54.7 

1.610 

165 

65.2 

1.825 

35 

21.4 

1.175 

79 

40.8 

1.395 

123 

55.0 

1.615 

166 

65.5 

1  830 

36 

22.0 

.180  ! 

80 

41.2 

1.400 

124 

55.2 

1620 

167 

65.7 

1.835 

37 

.22.5 

.185 

81 

41.6 

1.405 

125 

55.5 

1.625 

168 

65.9 

1.840 

38 

23.0 

.190  ; 

82 

420 

1.410 

126 

55.8 

1.630 

169 

66  1 

1.845 

39 

23.5 

.195 

83 

42.3 

1.415 

127 

56.0 

1.635 

170 

663 

1.850 

40 

24.0 

.200 

84 

42.7 

1.420 

128 

56.3 

1.640 

171 

66.5 

1.855 

41 

24.5 

.205 

85 

43.1 

1.425 

129 

56.6 

1.645 

172 

66.7 

1.860 

42 

25.0 

210 

86 

43.4 

1.430 

130 

56.9 

1.650 

173 

67.0 

1.865 

43 

25.5 

.215! 

87 

43.8 

1.435 

The  scale  of  the  Banme  hydrometer,  for  liquids  lighter  than  water,  is  so  adjusted 
that  the  point  to  which  the  spindle  \vould  sink,  in  a  solution  of  one  part  of  sodium 
chloride  in  nine  parts  of  water,  is  marked  0,  while  the  corresponding  point,  when 
pure  water  is  used,  is  marked  10.  The  instrument  used  for  liquids  heavier  than 
water  has  the  0  at  the  point  to  which  it  sinks  in  pure  water,  and  the  10  at  the  point 
reached  in  a  10  per  cent  solution  of  sodium  chlorine  at  17.5°. 

The  Twaddell  hydrometer  contains  two  hundred  divisions  for  a  difference  in 
specific  gravity  of  one.  The  0  of  Twaddell  is  the  specific  gravity  of  pure  water  and 
=  1.000.  The  instrument  is  usually  divided  into  five  sections  for  convenience.  The 
200  mark  would  correspond  to  a  specific  gravity  of  2.000. 


INDEX. 


Acetone,  360 

Acetylene,  368,  371,  374 

Acid,  acetic,  360  ;  antimonic,  343  ; 
arsenic,  336;  arsenic,  experiments 
with,  802;  arsenious,  337;  boric, 
354;  bromic,  166;  carbonic,  386  ; 
chlorauric,  614;  cbloric,  115  ; 
chloric,  experiments  with,  772  ; 
chlorous,  118;  chlorplatinic,  730  ; 
cblorsulphuric,  241  ;  chromic, 
652,  663  ;  chromic,  experiments 
with,  820;  cyanauric,  615  ;  cyan- 
aurous,  615  ;  cyanic,  404;  cyan- 
platinous,  730;  dichromic,  658  ; 
disilicic,  422;  disulphuric,  219; 
dithionic,  208,  225;  dititanic,  424; 
diuranic,  677;  ferric,  710;  ferrihy- 
drocyanic,  704  ;  ferrohydrocy- 
anic,  703  ;  fluoboric,  353;  fluo- 
silicic,  415;fluosilicic,  constitution 
of,  416 ;  fluosilicic,  preparation 
of,  811;  fluotantalic,  350;  fluo- 
titanic,  417;  fluozirconic,  417  ; 
formic,  380,  391 ;  fuming,  sulphu- 
ric, 219;  hydrobromic,  163;  hydro- 
bromic,  experiments  with,  776 ; 
hydrochloric,  43,  105;  hydrochlo- 
ric, experiments  with,  771  ;  hy- 
drochloric, formation  of,  770  ; 
hydrochloric,  preparation  of,  771; 
hydriodic,  169  ;  hydriodic,  reduc- 
ing action  of,  779;  hydrocyanic, 
402;  hydrofluoric,  178;  hydrofluo- 
ric, constitution  of,  179  ;  hydro- 
fluoric, experiments  with,  779  ; 
hydrosulphurous,  222;  hypobro- 
mous,  166;  hypochlorous,  116  ; 
hyponitrous,  282  ;  hypophos- 
phoric,  333;  hypophosphorous, 
333  ;  hyposulphurous,  208,  221, 
222;  iodic,171;  iodic,  constitution 
of,  176;  iodic,  experiments  with, 
779 ;  manganic,  182,  685;  meta- 
boric,  355;  metachromous,  659  ;^ 
metantimonic,  296,  343  ;  meta-* 
phosphoric,  296,  331;  metarsenic, 
296,  336  ;  metastannic,  643  ; 
rnetavanadic,  350;  molybdic,  671  ; 
muriatic,  43;  nitric,  43,  277  ; 


nitric,  experiments  with,  792  ; 
nitric,  normal,  263;  nitric,  prepa- 
ration of,  791;  nitric,  red  fuming, 
281;  nitric,  reduction  of,  794  ; 
nitrohydrochloric,  281  ;  nitrosyl- 
sulphuric,  212,  293  ;  nitrous,  281; 
nitrous,  experiments  with,  794  ; 
nitrous,  normal,  263;  Nordhausen 
sulphuric,  219;  normal  sulphuric, 
210,  218;  orthoantimonic,  343  ; 
orthophosphoric,296,326 ;  orthova- 
nadic,  350;  osmic,  724  ;  pentathi- 
onic,  208,  226  ;  perbromic,  166  ; 
perchloric,  118  ;  perchloric,  prep- 
aration of,  773  ;  periodic,  173  ; 
periodic,  constitution  of,  174  ; 
periodic,  normal,  174  ;  perman- 
ganic, 182,  687;  phosphoric,  326; 
phosphoric,  experiments  with, 
802;  phosphoric, glacial, 331;  phos- 
phoric, "insoluble,"  541;  phos- 
phoric, normal,  296 ;  phosphoric, 
"reverted,"  541;  phosphoric  "sol- 
uble, "  541 ;  phospho-molybdic,672; 
phosphorus,  332;  platinic,  729  ; 
prussic,  402  ;  pyroantimonic,  297, 
343;  pyroarsenic,  297,  336  ;  pyro- 
carbonic,  388;pyroligneous,  360  ; 
pyrophosphoric,  297,  330  ;  pyro- 
sulpharsenic,  342  ;  pyrosulpharse- 
nious,  342  ;  pyrosulphuric,  219  ; 
pyrotantalic,  351 ;  pyrovanadic, 
350;  ruthenious,  723;selenic,  242; 
selenious,  241;  silicic,  420;  silicic, 
experiments  with,  812  ;  silicic, 
insoluble,  422;  silicic,  normal,  420; 
stannic,  642  ;  sulphantimonious, 
346;  sulpharsenious,  340;  sulpho- 
carbonic,  406;  sulphocyamc,  407; 
sulphotelluric,  245  ;  sulphoxyar- 
senic,  342;  sulphuric,  43,  209;  sul- 
phuric, experiments  with,  782; 
sulphuric,  illustration  of  prepa- 
tion,  782;  sulphuric,  solid,  219, 
236;  sulphurous,  220;  sulphur- 
ous, experiments  with,  783;  sul- 
phydric,  200;  telluric,  244;  tellu- 
rious,  244;  tetrabortc,  354;  tetra- 
hydroxyl-sulphuric,  218;  tetrathi- 
843 


844 


INDEX. 


onic,  208,  226;  thiocarbonic,  406; 
thiosulphuric,  208,  223;  triazoic, 
277 ;  trithionic,  208,  225;  tungstic, 
674;  vanadic,  350;  zirconic,  424 

Acids,  61,  127;  avidity  of,  440;  basi- 
city of,  136;  constitution  of,  134; 
dibasic,  138  ;  monobasic,  138;  no- 
menclature of,  143;  organic,  380: 
pentabasic,  138  ;  polymolybdic, 
672;  polysilicic,  422;  polystannic, 
643;  polytungstic,  674;  silico- 
tungstic,  674;  tetrabasic,  138;  tri- 
bisic,  138;  trisilicic,  423 

Acids  of  antimony  and  arsenic,  con- 
stitution of,  346 

Acids  of  phosphorus,  constitution 
of,  335 

Acids  of  sulphur,  constitution  of, 
226 

Acidum,  phosphoricum,  glaciale, 
331 

Affinity,  chemical,  11,  435 

Affinity,  specific  coefficient  of,  442 

Affinity,  thermochemical  study  of, 
438 

Agate,  418 

Air,  248,  251 

Air,  analysis  of,  253,  785 

Alabandite,  684 

Alabaster,  539 

Albite,  575 

Alchemy,  4 

Alcohol,  379 

Aldehyde,  formic,  379 

Algaroth,  powder  of,  318,  347 

Alkalies,  127 

Alkali  metals,  482 

Alkaline  earths,  527 

Allanite,  655 

Allotropy,  89 

Alloys,  approximate  composition 
of,  839 

Allylene,  371 

Alum,  basic,  578;  burnt,  578;  insolu- 
ble, 579;  shale,  578 

Alumina,  aluminium  oxide,  575 

Alurninates,  571 

Aluminium,  563;  bronze,  566,  587; 
chloride,  566;  preparation  of,  818; 
hydroxide,  569;  oxide,  575  ;  sili- 
cates, 579;  sulphate,  576  ;  sul- 
phates, basic,  576 

Alums,  576 

Alunite,  578 

Amalgamation  process  for  silver, 
602 

Amalgams,  625 
Amethyst,  418 

Ammonia,  260,  266;  composition  of, 


270;  determination  of  composi- 
tion of,  790;  experiments  with, 
789;  preparation  of,  788 

Ammoniacal  liquor,  266,  365 

Ammonia-process  for  soda,  512 

Ammonium  alum,  579;  amalgam, 
273,  625;  carbamate,  522;  carbon- 
ate, 522  ;  chloride,  266,  519;  com- 
pounds, metallic  derivatives  of, 
274;  compounds,  structure  of,  274; 
hydrosulphide,  521;  hydroxide, 
270;  nitrate,  522;  nitrite,  267;  mag- 
nesium phosphate,  329,  560; 
phospho-molybdate,  672  ;  salts, 
269,  517;  salts,  formation  of,  789; 
sulphide,  520;  sulphide,  prepara- 
tion of,  817;  sulphocyanate,  519; 
uranate,  677 

Ammonium  sulphide  group,  198 

Ampere's  law,  73 

Anatase,  412,  423 

Anhydrides,  172,  209 

Anhydrite,  539 

Annealing,  544 

Antimony,  308;  and  its  compounds, 
experiments  with,  799;  blende, 
345;  oxy chlorides,  319,  347;  oxy- 
chloride,  experiment  with,  802; 
pentachloride,  318  ;  pentasul- 
phide,346;  pentoxide,  345;  salts  of, 
343;  sulphides,  experiments  with, 
802  ;  tetroxide,  345;  trichloride, 
318;  trioxide,  343;  trisulphide, 
345 

Antimony,  constitution  of  acids  of, 
346 

Antimonyl  salts,  344  ;  sulphate, 
344 

Apatite,  298,  529,  540 

Aqua  regia,  281 

Aragonite,  536 

Arbor  Saturni,  647 

Argentan,  716 

Argentic  nitrate,  608 

Argentous  chloride,  606 

Argon,  259 

Aristotle,  four  elements  of,  8 

Arsenic,  305  ;  and  its  compounds, 
experiments  with,798;  disulphide, 
340;  pentasulphide,  342  ;  pentox- 
ide, 340  ;  sulphides,  experiments 
with,  802  ;  trichloride,  317;  triox- 
ide, 338  ;  trioxide,  experiments 
with,  802;  trisulphide,  340 

Arsenic,  constitution  of  acids  of, 
346 

Arsenides,  305 

Arseniuretted  hydrogen,  306 

Arsenopyrite,  711 


INDEX. 


845 


Arsine,  306,  798 
Arsonium  compounds,  308 
Atomic  theory,  68;  weights,  71,  77, 

79,  824;  weights  and  specific  heat, 

449 

Atoms,  68,  74 
Aurates,  614 
Auric    chloride,    613;    compounds, 

610;  hydroxide,  614;  oxide,  614 
Aurous   chloride,  613;  compounds, 

610;  oxide,  614 
Avidity  of  acids,  440 
Avogadro's  law,  73 

Balance,  5 

Banca  tin,  640 

Barff's  process,  699 

Barite,  547 

Barium,  547;  carbonate,  549;  chlo- 
ride, 547;  chromate,  667;  dioxide, 
547;  hydroxide,  547;  nitrate,  549; 
oxide,  547;  peroxide,  547;  phos- 
phates, 550;  platinate,  724;  sul- 
phate, 549;  sulphide,  548 

Bases,  61,  127,  131;  acidity  of,  138; 
constitution  of,  134;  diacid,  138; 
monacid,  138;  nomenclature  of, 
144;  triacid,  138 

Basic-lining  process,  698 

Bauxite,  564,  571 

Bell-metal,  591 

Benzene,  372 

Berthollet,  15 

Beryll,  551 

Beryllium,  see  Glucinum. 

Bessemer  process,  698 

Bismuth,  311;  basic  nitrates  of ,  348, 
803;  dichloride,  319;  dioxide,  348; 
experiment  with,  799;  hydroxide, 
347;  nitrate,  311;  oxychloride, 
319,  349;  pentoxide,  349;  sub- 
nitrate,  348;  sulphate,  311;  tri- 
chloride, 319;  trioxide,  347;  tri- 
sulphide..  349 

Bisrnuthyl  salts,  348 

Black  ash,  511 

Black-lead,  359 

Bleaching  by  chlorine,  101,  103;  by 
sulphur  dioxide,  235;  Dowder, 
116,  534 

Block-tin,  640 

Blow-pipe,  398;  experiments  with, 
808 

Blue  vitriol,  595 

Bog  iron-ore,  706 

Boiling-points,  table  of,  826 

Bone-black,  362;  filters,  363;  ex- 
periments with,  803 

Boracite,  354 


Borax,  352,  515 

Boron,  351;  experiments  with,  803 
nitride,  356;  phosphate,  356;  salts 
of,  356;  trichloride,  352;  trifluo- 
ride,  356 

Boryl  potassium  tartrate,  356 

Brass,  591,  617 

Braunite,  679,  682 

Bricks,  583 

Brimstone,  crude,  188;  roll,  188 

Britannia  metal,  309,  640 

Bromides,  experiments  with,  812 

Bromine,  161;  Chloride,  166;  hy- 
drate, 162;  preparation  of,  776; 
water,  162 

Bromoform,  375 

Bronze,  591 

Brookite,  412,  423 

Brown  iron  ore,  694,  706 

Bunsen  burner,  400 

Burning,  32 

Butane,  369 

Butter  of  antimony,  318 

Butylene,  371 

Biityrum  antimonii,  318 

Cadmium,  622  ;  carbonate,  616  ; 
chloride,  622;  cyanide,  623;  sul- 
phate, 622;  sulphide,  622 

Caesium,  502 

Calamine,  617 

Calcium,  529;  carbide,  545;  carbon- 
ate, 536;  chloride,  530;  chloride, 
preparation  of,  817;  fluoride,  531; 
hydroxide,  532 ;  nitride,  545 ; 
oxide,  531;  phosphates,  540;  sili- 
cate, 542;  sulphate,  538;  sulphide, 
545 

Calc-spar,  536 

Calomel,  626 

Calorie,  37 

Calorimeter,  37 

Candle,  standard,  395 

Carbides,  368 

Carbon,  357  ;  amorphous,  360 ; 
chemical  conduct  of,  366;  diox- 
ide, 380  ;  dioxide,  experiments 
with,  805 ;  dioxide,  relation  to 
life,  385;  disulphide,  404;  experi- 
ments on  reducing  power  of,  804; 
monoxide,  389;  experiments  with, 
806;  tetrachloride,  375 

Carbonates,  386,  478 ;  acid,  387 ; 
basic,  387;  experiments  with,  815 

Carbonyl  chloride,  392 

Carbon  silicide,  424 

Carborundum,  424 

Carnallite,  483,  485 

Carnelian,  418 


846 


INDEX. 


Cassiterite,  639 

Cast-iron.  696;  gray,  696;  white,  696 

Cast-steel,  698 

Caustic  soda,  506 

Cavendish,  40 

Celestite,  546 

Cements,  545;  hydraulic,  545 

Cerite,  562,  655 

Cerium,  413,  655 

Cerussite,  646 

Chalcedony,  418 

Chalcocite,  589 

Chalcopyrite,  708 

Chalk,  529,  536 

Chance  process  for  recovery  of  sul- 
phur, 512 

Charcoal,  360;  absorption  of  gases 
by,  803;  animal,  362;  filters,  363; 
kiln,  361 

Chemical  action,  10,  426;  examples 
of,  737 

Chemical  change,  2 ;  caused  by 
electric  current,  735;  by  heat,  734 

Chemical  reactions,  causes  of,  434; 
kinds  of,  431;  ideal,  435 

Chemistry,  3;  organic,  358 

Chili  saltpetre,  507 

Chlorates,  476 

Chloraurates,  614 

Chlorides,  99,  460 ;  double,  465  ; 
experiments  with,  812  ;  general 
properties  of,  463;  of  acids,  209 

Chlorine,  96;  dioxide,  121;  experi- 
ments with,  769;  preparation  of, 
768;  hydrate,  104;  hydrate,  pre- 
paration  of,  770  ;  liquid,  104 ; 
monoxide,  120;  trioxide,  121 

Chlormercurates,  622 

Chloro-acids,  465 

Chloroaluminates,  568 

Chlorostannates,  642 

Chloroform,  375 

Chlorophyll,  695 

Choke-damp,  384 

Chlorplatinates,  730 

Chromates,  663;  experiments  with, 
820 

Chrome-alums,  662 

Chrome-red,  668 

Chrome-yellow,  668 

Chromic  chloride,  660;  compounds, 
658;  hydroxide,  661;  iron,  659; 
oxide,  662;  sulphate,  662 

Chromite,  573,  657,  659 

Cromium,  657;  oxychloride,  668; 
trioxide,  666 

Chromous  chloride,  660;  com- 
pounds, 658;  hydroxide,  661 

Chromyl  chloride,  668 


Chrysoberyl,  573 

Cinnabar,  624,  631 

Clay,  581 

Cleveite,  585 

Coal,  364;  anthracite,  364;  bitumi- 
nous, 364;  gas,  394;  experiments 
with,  807;  oil,  358;  tar,  365 

Cobalt,  713;  cyanides,  714;  sulphide, 
714 

Cobaltic  hydroxide,  714;  oxide,  714 

Cobaltite,  713 

Cobaltous  chloride/  713;  cobaltic 
oxide,  714;  hydroxide,  714;  oxide 
714 

Coke,  362,  394 

Columbite,  350 

Columbium,  351 

Combination,  11,  24;  direct,  431 

Combining  weights,  17,  71 

Combustion,  34;  experiments  on, 
749;  Lavoisier's  explanation  of,  34 

Compound,  chemical,  10,  12,  736 

Condy's  liquid,  688 

Constitution,  80 

Copper,  587;  basic  carbonate,  387  ; 
hydride,  592;  metallurgy  of,  589  ; 
native,  589;  peroxide,  595;  plat- 
ing, 600;  ruby,  589;  suboxide,  595 

Copperas.  709 

Corrosive  sublimate,  627 

Corundum,  575 

Cotunnite,  648 

Crocoisite,  646,  659 

Cryolite,  177,  466,  503;  constitution 
of,  180 

Cupric  arsenite,  598;  carbonates, 
598;  chloride,  592;  cyanide,  599; 
hydride,  592;  hydroxide,  594;  ni- 
trate, 598;  oxide,  594;  sulphate, 
595;  sulphide,  600;  sulphocyanate, 
599 

Cuprite,  594 

Cuprous  chloride,  592;  chloride, 
preparation  of,  818;  cyanide,  599; 
hydride,  592;  hydroxide,  593; 
iodide,  593;  oxide  594;  sulphide, 
599;  sulphocyanate,  599 

Cyanides,  401,  407 

Cyanogen,  401;  experiments  with, 
810 

Cyanplatinites,  730 

Dal  ton,  15,  68 
Datholite,  354 
Davy,  96 
Deacon's  process, 
Decay,  381 

Decomposition,  11,  24;  direct,  432; 
double,  25 


INDEX. 


847 


Decrepitation,  505 

De  la  Bastie  glass,  544 

Deliquescence,  58 

Deliquescent  salts,  764 

Developers  in  photography,  607 

Dialyser,  421 

Dialysis,  421 

Diamond,  359 

Diaspore,  570 

Di-chlor-ethane,  373;  methane,   375 

Dichro-cobaltic  chloride,  715 

Dichromates,  experiments  with,  820 

Didirniuni,  351,  656 

Diffusion,  45;  experiments  on,  757 

Dimorphism,  190 

Di-sodiuin  pyroantimonate,  515 

Dissociation,  61,  314,  443;  of  dis- 
solved substance,  453 

Distillation,  67;  destructive,  360; 
dry,  360;  experiment  on,  764 

Dolomite,  380,  555 

Double  chloride,  465;  union,  372 

Double  salts,  319 

Drummond  light,  55 

Dumas,  his  study  of  combustion  of 
hydrogen,  49;  method  for  deter 
minating  the  specific  gravity  of 
vapors,  765 

Earthenware,  582 

Efflorescence,  58,  508 

Efflorescent  salts,  764 

Eka-aluminium,  635 

Eka-boron,  584 

Eka-silicon,  639 

Electrolysis,  443,  445 

Electrolytic  dissociation,  446 

Electrolytic  process  for  chlorine, 
99 

Electro-negative  ions,  446 

Electro-positive  ions,  446 

Electrotypes,  600 

Elements,  8,  19,  21;  acid-forming. 
112;  base-forming,  111,  455;  melt- 
ing-points and  boiling-points,  825 
of  Aristotle,  8  ;  replacing  power, 
83;  symbols  of,  19,  21 

Emerald,  552 

Emery,  575 

Energy,  chemical,  38;  conservation 
of,  6;  stored  up  in  plants,  385 

Enstatite,  422 

Epsom  salt,  559 

Equations,  chemical,  23 

Erbium,  562 

Etching  on  glass,  179 

Ethane,  369,  371 

EtMops  martialis,  631 

Ethylene,  371,  374 


Euchlorine,  121 
Eudiometer,  50 

Eudiometric    method,    50;    experi- 
ments, 763 
Euxenite,  562,  585 

Faraday,  96 

Feldspar,  480,  564 

Ferment,  nitrifying,  277 

Fermentation,  alcoholic,  381 

Ferric  chloride,  695 ;  chloride,  pre- 
paration of,  821;  ferrocyanide, 
703;  hydroxide,  706;  soluble, 
707;  oxide,  707;  sulphate,  710; 
sulphate,  preparation  of,  821; 
sulphide,  708 

Ferrous  ammonium  sulphate,  pre- 
paration of,  821;  carbonate,  708; 
chloride,  700;  ferricyanide,  704; 
hydroxide,  705;  oxide,  706;  phos- 
phate, 710;  sulphate,  708;  sul- 
phate, preparation  of,  821;  sul- 
phide, 707;  ferric  oxide,  707 

Fire-damp,  373 

Flame,  395;  oxidizing,  397;  reac- 
tions, 525;  reducing,  397 

Flames,  395;  causes  of  luminosity 
of,  399;  structure  of,  397 

Flint,  418 

Flores  zmci,  620 

Fluocolumbates   351 

Fluorides,  constitution  of,  179 

Fluorine,  177 

Fluor-spar,  177,  529 

Fluosilicates,  415 

Fluostannates,  642 

Fluotantalates,  350 

Fluotitanates,  417 

Flux,  531 

Fool's  gold,  711 

Formulas,  constitutional,  81;  molec- 
ular, 79 

Franklinite,  617,  706 

Freezing-mixtures,  840 

Fulminating  mercury,  632 

Fumaroles,  354 

Gadolinite,  562,  584,  655 

Gadolinium,  585 

Gahnite,  573,  617 

Galenite,  646 

Gallium,  635 

Gallic  sulphate,  635 

Gallons  chloride,  635 

Garnet,  542 

Gases,    kindling    temperature    of, 

807;  measurement  of  volume  of, 

741;  weights  of,  839 
Gay  Lussac  tower,  214 


848 


INDEX. 


Germanium,    425,    639  ;    chloride, 

425;  oxide,  425 
German  silver,  591,  716 
Gersdorffite,  716 
Glass,  542 
Glauber,  105 
Glauber's  salt,  507 
Glover  tower,  214 
Glucinum,    551;     carbonate,     554; 

chloride,  552;  hydroxide,  553 
Gold,   610;    alloys,   613;  amalgam, 

626;  dichloride,  613;   metallurgy 

of,  611;  sulphide,  615 
Graham,  47 
Graphite,  359 
Greenockite,  622 
Green  Vitriol,  709 
Guignet's  green,  662 
Gun-metal,  591 
Gunpowder,  493 
Gypsum,  529,  538 

Halogens,  160 

Hardness,  permanent,  538;  tem- 
pory,  538 

Hard  water,  538 

Hauerite,  681 

Hausmannite,  679,  681 

Heat  of  combustion,  36;  of  decom- 
position, 38;  of  neutralization, 
440;  specific  law  of,  449 

Heavy  spar,  547 

Helium.  585 

Hematite,  694 

Hepar  sulfur  is,  491 

Heptane,  369 

Hexane,  369 

Holmium,  562 

Homologous  series,  370 

Homology,  370 

Hornblende,  555 

Hydrargillite,  569 

Hydrates,  61,  144,  146 

Hydraulic  cement,  545;  mining,  612 

Hydrazine,  275 

Hydrocarbons,  368,  804 

Hydrochloric  acid  hydrate,  109 

Hydrogen,  40 ;  amount  evolved 
when  a  known  weight  of  a  metal 
is  dissolved  in  an  acid,  754;  burn- 
ing of,  49,  106,  759;  chemical 
properties  of,  758;  dioxide,  91; 
dioxide,  experiments  with,  768; 
peroxide,  91;  persulphide,  201; 
preparation,  751;  selenide,  204; 
sulpliide,  193;  sulphide,  experi- 
ments with,  780;  sulphide  group, 
198;  sulphide  in  chemical  analy- 
sis, 196;  telluride,  205 


Hydrogenium,  47 
Hydrogen- valence,  156 
Hydrometer  table,  841 
Hydrosulphides,  200,  474 
Hydroxides,   61,   469;    experiments 

with,  814;  formation  from  oxides, 

236 

Hydroxol,  237 
Hydroxylamine,  265,  276 
Hypochlorites,  476 
Hyposulphite  of  soda,  508 

Iceland  spar,  536 

Ice-machine,  268 

Illuminating  gas,  394 

Illumination,  394 

Indium,  636 

Infusorial  earth,  420 

Ink,  sympathetic,  714 

lodic  anhydride,  172 

Iodides,  experiments  with,  812 

Iodine,  167;  bromide,  177;  chloride, 

176;  experiments  with,  778;  pen- 

tafluoride,    180;    pentoxide,    172; 

trichloride,  177 
lodoform,  375 
Ions,  446 
Iridium,  724;  chlorides,  725;  oxides, 

725 
Iron,    691;    carbonyls,    711;     disul- 

phide,     711;     galvanized,      618; 

metallurgy,  695 
Isomerism,  ^30 

Kainite,  499,  555 
Koaline,  564,  580 
Kelp,  167 
Kieserite,  555 

Kindling  temperature,  34;  of  gases, 
395 

Lamp  black,  362 

Lanthanum,  655 

Lapis-lazuli,  581 

Laughing  gas,  284 

Lavoisier's  work,  5 

Law  of  definite  proportions,  14;  of 
indestructibility  of  matter,  6;  of 
multiple  proportions,  15 

Lead,  425,  646;  its  com  pounds  r 
experiments  with,  820;  carbonate, 
652;  chloride,  648;  chromate,  668; 
hydroxide,  649;  iodide,  648; 
molybdate,  672;  metallurgy  of, 
646;  nitrate,  652;  oxide,  649;  pen- 
cils,  359;  peroxide,  650;  sesqui- 
oxide,  650;  suboxide,  649;  sugar 
of,  650;  sulphate,  654;  sulphide, 
652;  tree,  647 


INDEX. 


849 


Le  Blanc  process  for  soda,  510 

Lepidolite,  516 

Lignite,  364 

Lime,  531;  chloride  of,  116;  slaked, 
531 

Liine-light,  55 

Lime-stone,  529,  536 

Lime-water,  532 

Linnaeite,  714 

Liquor  Ferri  chlorati,  701;  Ferri 
sesquichlorati,  701;  Stibii  muria- 
tici,  319 

Litharge,  649 

Lithium,  516;  carbonate,  517;  chlor- 
ide, 517;  phosphate,  517 

Liver  of  sulphur,  491 

Loadstone,  707 

Lunar  caustic,  609 

Luteo-cobaltic  chloride,  715 

Magnesia,  558;  alba,  560;  usta,  558 

Magnesio  ferrite,  573 

Magnesite,  555 

Magnesium,  555;.  borate,  561;  car- 
bonate, 559;  chloride,  557;  chlor- 
ide, preparation  of,  817;  hydrox- 
ide, 558;  phosphates,  560;  prepa- 
ration of,  817;  oxide,  558;  silicate, 
561;  silicide,  561;  sulphate,  559 

Magnetite,  694,  707 

Malachite,  589,  598 

Malaria,  259 

Manganates,  685 

Manganese,  182,  678;  black  oxide  of , 
682;  dioxide,  682;  hept-oxide,  689; 
salts,  preparation  of,  821;  tetra- 
chloride,  680;  tetrafluoride,  680; 
trichloride,  680 

Manganic  oxide,  682;  sulphate,  685 

Manganite,  679,  682 

Manganites,  683 

Manganous  carbonate,  685;  chlo- 
rides, 679;  cyanide,  685;  hydrox- 
ide, 681 ;  maganic  oxide,  681 ;  ox- 
ide, 681;  sulphate,  685;  sulphide, 
684 

Marble,  529,  536 

Marcasite,  711 

Marl.  537,  581 

Marsh  gas,  368,  369,  373 

Marsh's  test'  for  arsenic,  310,  797 

Mass  action,  436 

Mass,  influence  of,  436 

Matches,  302;  safety,  302 

Matter,  constitution  of,  68;  early 
views  regarding  composition  of, 
7;  law  of  indestructibility  of,  6 

Meerscbaum,  555 

Melting-point,  table  of,  826 


Mendelejeff,  148 

Mercuric  chloramide,  633;  chloride, 
627;  compounds,  623;  cyanide, 
631;  diammonium  chloride,  634; 
iodide,  629;  nitrate,  633;  nitrate, 
preparation  of,  819;  oxide,  630; 
sulphide,  631 

Mercurias  solubilis  HaJmemanm, 
634 

Mercurous  ammonium  chloride,  633; 
chloramide,  633;  chloride,  626; 
compounds,  623;  iodide,  628; 
nitrate,  632;  nitrate;  preparation 
of,  819;  oxide,  630;  sulphide,  630 

Mercury,  623;  metallurgy  of,  624 

Metallic  properties,  456 

Metallurgy,  458 

Metals,  111,  134,  456;  properties, 
459 

Metastannates,  643 

Metathesis,  25,  433 

Metatungstate,  674 

Metavanadates,  350 

Meteorites,  694 

Methane,  369,  373 

Methyl  alcohol,  360,  379 

Meyer,  Lothar,  148 

Meyer,  Victor,  method  for  deter- 
mining the  specific  gravity  of  va- 
pors, 766 

Mica,  542;  564 

Microcosmic  salt,  332;  523 

Minerals,  457 

Miner's  safety  lamp,  396 

Minium,  651 

Mixtures,  mechanical,  12,  736 

Molecular  weights,  75,  77 

Molecules,  73,  74,  427 

Molybdates,  671 

Molybdenite,  670 

Molybdenum,  670;  chlorides,  670; 
oxides,  671 

Mono-chlor-methane,  375 

Mortar,  544 

Mosaic  gold,  644 

Nascent  state,  90 
Neodymium,  351,  656 
Neutralization,     127;     experiments 

on,  774 

Newlands,  148 
Newton's  metal,  312 
Niccolite,  716 
Nickel,  716;   alloys  ,716;    carbonyl, 

717;  cyanide,  717;  plating,  717 
Nickelic  hydroxide,  717 
Nickelous  chloride,  717,  hydroxide, 

717;  oxide,  717 
Niobium,  350 


850 


INDEX. 


Nitrates,  475;  basic,  280;  formation 
of,  793;  formation  of,  in  the  soil, 
261 

Nitric  oxide,  285;  experiments  with, 
795 

Nitrification,  277,  492 

Nitre-compounds,  280 

Nitrogen,  248;  boride,  356;  experi- 
ments with,  785;  iodides,  292; 
pentoxide,  289;  peroxide,  288; 
peroxide,  experiments  with,  795; 
preparation  of,  784;  relations  to 
life,  257;  structure  of  compounds 
of,  289;  sulphide,  293;  trichlo- 
ride, 292;  triiodide,  292;trioxide, 
287;  trioxide,  experiments  with, 
795 

Nitroprussiates,  705 

Nitrosyl  chloride,  281 

Nitrous  oxide,  283;  experiments 
with,  794 

Non-metals,  112,  456 

Octane,  369 

Olefiant  gas,  374 

Olivine,  561 

Opal,  418 

Ores,  458 

Orpiment,  305,  340 

Orthite,  562 

Orthoclase,  423 

Osmium,  723;  chlorides,  724;  oxides, 
724 

Oxidation,  slow,  35 

Oxides,  39,  467;  acidic,  172,  209; 
basic,  172 

Oxygen,  28;  amount  liberated  from 
a  known  weight  of  potassium 
chlorate,  745;  and  acid  properties, 
141;  burning  of,  807;  chemical 
properties  of ,  747;  physical  prop- 
erties of,  747;  preparation  of,  740; 
valence,  156 

Oxyhydrogen  light,  55;  blow-pipe, 
54;  blow-pipe,  experiments  with, 
763 

Oxysulphides,  406 

Ozone,  85;  experiments  with,  767; 
in  the  air,  88 

Palladic  chloride,  726;  oxide,  727 

Palladious  chloride,  726;  oxide,  727 

Palladium,  726;  hydrogen,  726; 
suboxide,  727 

Paracyanogen,  632 

Passive  state  of  iron,  700 

Pattison's  method  of  separating  sil- 
ver from  lead,  602 

Peat,  364 


Pentane,  369 

Periodates,  174 

Periodic  law,  148,  158,  429 

Permanent  white,  549 

Permanganates,  687 

Petalite,  516 

Petroleum,  358,  369 

Philosopher's  wool,  819 

Phlogiston  theory,  33 

Phosgene,  392 

Phosphates,  479;  acid,  329;  neutral, 
329;  normal,  329;  primary,  329; 
secondary,  329;  tertiary;  329 

Phosphine,  302;  experiments  with, 
797;  liquid,  304 

Phosphonium  bromide,  304;  chlo- 
ride, 304;  iodide,  304 

Phosphoric  anhydride,  334 

Phosphorite,  298,  529,  540 

Phosphorus,  298;  anhydride,  334; 
crystallized,  metallic,  301 ;  exper- 
iments with,  796;  iodides,  317; 
oxychlorides,  336;  pentachloride, 
3J4;  pentachloride,  action  on  hy- 
droxides, 316;  pentachloride, 
preparation  of,  801;  pentafluor- 
ide,  317;  pentoxide,  334;  red,  301; 
suboxide,  335;  tetroxide,  335;  tri- 
bromide,  317;  trichloride,  812; 
trichloride,  preparation  of,  799; 
trioxide,  334 

Phosphorus,  constitution  of  acids 
of,  335 

Phosphuretted  hydrogen,  302 

Photography,  607 

Photometerj  395 

Physical  change,  2 

Physics,  3 

Pig-iron,  696 

Pinchbeck,  591 

Pink  salt,  642 

Pithcblende,  675 

Placer  mining,  611 

Plaster  of  Paris,  539 

Platinic  chloride,  729;  chloride, 
preparation  of,  821;  hydrox:de, 
730;  oxide,  730;  sulphide,  730 

Platinous  chloride,  729;  hydroxide, 
730;  oxide,  730;  sulphide,  730 

Platinum,  727;  alloys,  728;  bases, 
731;  black,  727;  cyanides,  729; 
metals,  722;  spongy,  727 

Plumbago,  359 

Plurnbates,  651 

Plumbites,  650 

Polymerism,  515 

Polysulphides,  207 

Porcelain,  582 

Potash,  483,  500 


INDEX. 


851 


Potassium,  483;  alum,  578;  bro- 
mide, 485;  carbonate,  acid,  500; 
carbonate,  extraition  from,  wood 
ashes,  816;  chlorate,  114,  494; 
chlorate,  preparation  of,  772; 
chloride,  485;  chloriridate,  725; 
chlorthorate,  417;  chrornate,  664; 
cyan  ate,  498;  cyanate,  preparation 
of,  810;  cyanide,  496;  cyanide, 
preparation  of,  810;  dichromate, 
664;  diperiodate,  496;  disulphite, 
500;  ferricyanide,  704;  ferrocya- 
nide,  401,  702;  fluogermanate,639; 
fluoride, 485;  fluosilicate,115;  fluo- 
thorate,  417;  fluotitanate,  412; 
hydride,  485;  hydrosulphide,  490; 
hydroxide,  488;  hydroxide,  prep- 
aration of,  818;  hypochlorite,  114; 
iodide,  485;  iodide,  preparation  of, 
816;  manganate,  686;  manganate, 
preparation  of,  821;  rnesoperio- 
date,  496;  nitrate,  492;  nitrite, 
494;  osinite,  724;  oxide,  489; 
perchlorate,  118,  495;  perchlorate, 
preparation  of,  773;  periodate, 
496;  permanganate,  183,  687;  per- 
inangate,  preparation  of,  821;  per- 
oxide, 489  ;  perruthenite,  723; 
phosphates,  500 ;  polysulphides, 
491;  ruthenite,  723;  silicate,  501; 
stannate,  643;  sulphantimonate, 
492;  sulpharsenate,  492;  sulph- 
arsenite,  492;  sulphate,  498;  sul- 
phate, acid,  499;  sulphide,  490; 
sulphite,  500;  sulphite,  acid,  500; 
sulphocyanate,  498;  tetrachro- 
mate,  668;  tichromate,  668 

Praseo-cobaltic  chloride,  715 

Praseodymium,  351,  656 

Preparing  salt,  643 

Priestley,  28 

Printing-ink,  362 

Propane,  369 

Propylene,  371 

Proust,  15 

Prussian  blue,  704 

Puddling,  697 

Purple  of  Cassius,  614 

Purpureo-cobaltic  chloride,  715 

Pyrargyrite,  603 

Pyrolusite,  679 

Pyrophospbates,  331 

Pyrosiderite,  706 

Pyrrhotite,  694 

Quartation.  613 
Quartz,  410,  418 
Quick-lime,  531 


Raoult's  method  for  determining 
molecular  weights,  521 

Reaction,  chemical,  24;  endother- 
niic,  94;  exothermic,  94 

Realgar,  305,  340 

Red  lead,  651 

Reducing  agent,  47 

Reduction,  47;  by  carbon,  368;  by 
hydrogen,  759 

Respiration,  383 

Reversion  of  phosphates,  541 

Rhodium,  724 

Rhodocroisite,  679 

Rinmann's  green,  622 

Roasting,  474 

Rock-crystal,  418 

Roseo-cobaltic  chloride,  715 

Rose's  metal,  312 

Rouge,  717 

Rubidium,  502 

Ruby,  575 

Ruby  copper,  594 

Rupert's  drops,  544 

Ruthenium,  722;  chloride,  722;  ox- 
ide, 723 

Rutile,  412,  423 

Safety  lamp,  374 

Sal  ammoniac,  266,  519 

Salt,  common,  504 

Saltpeter,  492;  plantations,  492 

Salts,  131,  139;  acid,  139;  basic,  138; 
constitution  of,  136;  decomposi- 
tion of  by  bases,  469;  formation 
of,  461;  neutral,  137;  nomen- 
clature of,  144;  normal,  137, 
139 

Samarium,  585 

Samarskite,  585 

Sand,  418 

Sapphire,  575 

Scandium,  584 

Scheele,  28 

Scheele's  green,  598 

Scheelite,  673 

Schlippe's  salt,  346,  506 

Schweinfurt  green,  598 

Selenium,  203;  acid  chlorides  of, 
243;  dioxide,  243 

Serpentine,  423,  555 

Siderite,  694,  702 

Siemens-Martin  furnace,  698 

Silicates,  480 

Silicides,  424 

Silicon,  410;  dioxide,  418;  hexa- 
chloride,  414;  hydride,  412;  mag- 
nesium, 557;  preparation  of,  810; 
tetrachloride,  413;  tetrafluoride, 


652 


INDEX. 


414;  tetrafluoride,  preparation  of, 
811 

Silver,  601;  allotropic  forms  of,  605; 
alloys  of,  605 ;  amalgam,  606  ; 
borates,  609;  bromide,  606;  chlo- 
ride, 606;  chromate,  668;  cyanide, 
609;  iodide,  606;  metallurgy  of, 
602;  nitrate,  608;  octoborate, 
609:  oxide,  608;  peroxide,  608; 
suboxide,  608;  sulphide,  608;  sul- 
phocyanate,  609;  tree,  prepara- 
tion of,  819;  triazoate,  608 

Slaking,  532 

Smalt,  715 

Smaltite,  713 

Smithsonite,  617 

Soapstone,  555 

Soda,  calcined  purified,  511;  crude, 
511;  crystallized,  511;  from  cryo- 
lite, 513 

Soda-water,  383 

Sodium,  502;  alum,  579;  amalgam, 
625;  ammonium  phosphate,  523; 
borate,  515  ;  bromide,  506  ;  car- 
bonate, 509;  carbonate,  acid  or 
primary,  514;  carbonate,  prepar- 
ation of,  by  ammonia  process, 
816  ;  chloride,  504 ;  chromate, 
667;  dichromate,  667;  fluoride, 
506 ;  hydride,  504  ;  hydroxide, 
506;  iodide,  506;  nitroprussiate, 
705  ;  metaphosphate,  329,  515  ; 
monoxide,  506;  nitrate,  506;  pan- 
tungstate,  674  ;  peroxide,  506  ; 
phosphates,  514;  potassium  car- 
bonate, 514;  pyrophosphate,  330; 
silicate,  516;  stannate,  643;  sul- 
phantimonate,  506 ;  sulphate, 
507 ;  sulphate,  experiment  on 
supersaturated  solution  of,  816; 
thiosulphate,  224,  508;  uranate, 
677 

Solder,  soft,  640 

Solubilities,  table  of,  826 

Solution,  62;  as  an  aid  to  chemical 
action,  64 

Solvay  process  for  soda,  512 

Spathic  iron,  694,  708 

Specific  heat  and  atomic  weights, 
449 

Specific  gravity  of  vapors,  deter- 
mination of,  765 

Spectroscope,  525 

Spectrum,  banded,  526;  continuous, 
526 

Sphalerite,  617 

Spiegel  iron,  697 

Spinels,  572 

Spiritus  fumans  Libavii,  642 


Spitting  of  silver,  604 

Spodumene,  516 

Stahl,  33 

Stalactites,  537 

Stalagmites,  537 

Stannic  chloride,  642 ;  chloride, 
preparation  of,  820;  hydroxide, 
642;  oxide,  644;  salts,  645;  sul- 
phate, 645;  sulphide,  644 

Stannite,  634 

Stannous  chloride,  641;  chloride, 
preparation  of,  819;  hydroxide, 
642;  oxide,  643;  salts,  645;  sul- 
phate, 645;  sulphide,  644 

Stas,  15 

Steel,  698 

Stibine,  309,  799 

Stibnite,  308,  345 

Stolzite,  673 

Strass,  543 

Stromeyerite,  608 

Stontianite,  546 

Strontium,  546 ;  chloride,  546  ; 
hydroxide,  546 ;  nitrate,  546  ; 
oxide,  546;  sulphate,  546 

Sublimation,  519 

Substitution,  41,  102 

Superphosphate  of  lime,  541 

Sulphantimonates,  346 

Sulphates,  476;  experiments  with, 
814 

Sulphides,  191,  472 

Sulphites,  2^2,  478;  acid,  222;  nor- 
mal,  222 

Sulpho-salts,  475 

Sulphostannates,  644 

Sulphur,  187 ;  acid  chlorides  of, 
239  ;  auraturu,  346  ;  dichloride, 
202  ;  dioxide,  233  ;  dioxide,  ex- 
periments with,  783;  experiments 
with,  779;  flowers  of,  188;  hept- 
oxide,  233;  hexiodide,  203;  in- 
soluble, 191;  monochloride,  202; 
sesquioxide,  233;  soluble,  191; 
stick,  188 ;  tetrachloride,  202 ; 
trioxide,  236;  trioxide,  experi- 
ments with,  784;  waters,  187 

Sulphur  auratum,  346 

Sulphur,constitution  of  acids  of  ,226 

Sulphuretted  hydrogen,  193 

Sulphuryl  chloride,  240 

Sulphuryl-hydroxyl  chloride,  241 

Sulphydrates,  200 

Sylvite,  483,  485 

Symbols  of  compounds,  22;  of  ele- 
ments, 19,  21 

Tachjdrite,  530 
Talc,'  555 


INDEX. 


853 


Tantalite,  350 

Tantalum,  350 

Tartar,  crude,  483;  emetic,  344 

Tellurium,  204;  dioxide,  245;  tetra- 
chloride,  205;  trioxide,  245 

Tempering  of  glass,  544;  of  steel, 
698 

Tetra-chlor-ethane,  373 

Tetra-chlor-metbane,  375 

Tballic  chloride,  636;  hydroxide, 
637;  sulphate,  637 

Thallium,  636 

Thallous  chloride,  636;  hydroxide, 
636;  phosphate,  637;  sulphate, 
637;  sulphide,  637 

Theory,  use  and  value  of  a,  70 

Thermocheinical  measurements, 
value  of,  439 

Thermochemistry,  27 

Thionyl  chloride,  239 

Thomas- Gilchrist  process,  698 

Thorite,  413 

Thorium,  413;  dioxide,  424;  tetra- 
chloride,  417;  tetrafluoride,  417 

Thulium,  562 

Tin,  425  ;  amalgam,  640  ;  metal- 
lurgy of,  639;  salt,  641;  stone, 
639,  644 

Titanium,  412;  dioxide,  423;  sul- 
phate, 424;  tetrachloride,  416; 
tetrafluoride,  417 

Titanyl  sulphate,  424 

Toluene,  372 

Tri-chlor-methane,  375 

Tridymite,  418 

Triple  linkage,  373;  union,  373 

Tungstates,  673 

Tungsten,  673;  chlorides,  673; 
oxides,  674 

Turnbull's  blue,  705 

Tuyeres,  695 

Type  metal,  309 

Ultramarine,  581 
Uranates,  677 
Uraninite,  675 

Uranium,  675;  chlorides,  676;    ox- 
ides, 676;  yellow,  677 
Uranous  salts,  676 
Uranyl  salts,  247,  676 

Valence,  81,  428;  variations  of,  155 
Vanadium,     350;     chlorides,    350; 
oxides,  350 


Vein-mining,  611 

Ventilation,  258 

Vitriol,  oil  of,  211 

Vitriols,  596 

Volumes,  law  of  combination  by,  45 

Water,  57;  as  a  solvent,  62;  con- 
tamination of,  by  sewage,  67 ; 
determination  of  composition  of, 
760;  in  organic  substances,  763; 
of  crystallization,  57,  763;  proofs 
of  composition  of,  58;  purifica- 
tion of,  67;  synthesis  of,  59; 
transformation  into  earth,  5 

Water-gas,  43,  389 

Water-glass,  501,  516 

Waters,  chalybeate,  66;  efferves- 
cent, 66;  natural,  65;  sulphur,  66 

Weights,  atomic,  19,  21;  combin- 
ing, 17, 18 

Weights,  length  and  volume,  com- 
parison of,  840 

Weldon's  process,  99,  682 

White  lead,  653 

White  precipitate,  633;  fusible, 
634;  infusible,  633 

Witherite,  547 

Wolframite,  673 

Wollastonite,  542 

Wood's  metal,  312 

Wood  spirits,  360 

Work,  chemical,  38 

Wrought  iron,  697 

Wulfenite,  646 

Xylene,  372 

Yellow  prussiate  of  potash,  401 
Ytterbium,  581 

Yttrium,  584  ;  chloride,  584 ; 
hydroxide,  584;  oxide,  584 

Zeolites,  542 

Zinc,  617;  alloys,  618;  blende,  614; 
carbonate,  621;  chloride,  618; 
dust,  617;  experiment  on  burning 
of,  819;  hydroxide,  619;  metal- 
lurgy of;  617;  oxide,  620;  solu- 
tion of,  in  sodium  hydroxide,  819; 
sulphate,  621  ;  sulphide,  620  ; 
white,  620 

Zircon,  413,  424 

Zirconium,  413;  dioxide,  424;  oxy- 
fluoride,  417;  tetrachloride,  417; 
tetrafluoride,  417 


SCIENCE 
REFERENCE  AND  TEXT-BOOKS 

PUBLISHED    BY 

HENRY  HOLT  &  COMPANY,  2? 


Books  marked  *  are  chiefly  for  reference  and  supplementary  use,  and  are 
to  be  found  in  Henry  Holt  &>  Co.'s  List  of  Works  in  General  Literature. 
For  further  particulars  about  books  not  so  marked  see  Henry  Holt  &*  CoSs 
Descriptive  Educational  Catalogue.  Excepting  JAMES'S  PSYCHOLOGIES, 
WALKER'S  POLITICAL  ECONOMIES,  and  ADAMS'  FINANCE,  all  in  the  Ameri- 
can Science  Series,  this  list  contains  no  works  in  Philosophy  or  Economics. 

American  Science  Series 


1.  Astronomy.     By  SIMON  NEWCOMB,  Professor  in  the  Johns  Hopkins  University, 

and  EDWARD  S.  HOLDEN,  late  Director  of  the  Lick  Observatory,  California. 

Advanced  Course.     512  pp.     8vo.     $2.00  net. 
The  same.     Briefer  Course.     352  pp.     12010.     $1.12  net. 
The  same.   Elementary  Course.    By  E.  S.  HOLDEN.   446pp.    i2mo.     %\.i.onet. 

2.  Zoology.    By.  A.  S.  PACKARD,  Jr.,  Professor  in  Brown  University.    Advanced 

Course.     722  pp.     8vo.     $2.40  net. 
The  same.     Briefer  Course.     338  pp.     $1.12  net. 
The  same.     Elementary  Course.     290  pp.     i2mo.     80  cents  net. 

3.  Boiany.     By   C.    E.    BESSEY,    Professor    in    the    University    of    Nebraska. 

Advanced  Course.     611  pp.     8vo.     $2.20  net. 
The  same.     Briefer  Course.     356  pp.     $1.12  net. 

4.  The  Human  Body.     By  H.  NEWELL  MARTIN,  sometime  Professor  in  the  Johns 

Hopkins  University. 

Advanced  Course.     685   pp.     8vo.     $2.50  net.     Copies  without  chapter  on 

Reproduction  sent  when  specially  ordered. 
The  same.    Briefer  Course.   (Entirely  neiu  edition  revised  by  Prof.  G.  Wells 

Fitz,  of  Harvard.)    408  pp.     i2ino.     $1.20  net. 
The  same.     Elementary  Course.     261  pp.     iamo.     75  cents  net. 
The  Human  Body  and  the  Effect  of  Narcotics.     261  pp.     i2mp.     $1.20  net. 

5.  Chemistry.     By    IRA   REMSEN,    Professor    in     Johns    Hopkins     University. 

Advanced  Course.     850  pp.     8vo.     $2.80  net. 
The  same.    Briefer  Course.     435  pp.    $1.12  net. 
The  same.     Elementary  Course.     272  pp.     12010.     80  cents  net. 
Laboratory  Manual  (to  Elementary  Course).     196  pp.     12010.     40  cents  net. 
Chemical  Experiments.     By  Prof.  REMSEN  and  Dr.  W.  W.  RANDALL.     (For 

Briefer  Course.)    No  blank  pages  for  notes.    158  pp.    i2mo.    50  cents  net. 

6.  Political  Economy.     By  FRANCIS  A.  WALKER,  President  Massachusetts  Insti- 

tute of  Technology.     Advanced  Course.     537  pp.     8vo.     $2.00  net. 
The  same.    Briefer  Course.    415  pp.     i2mo.     $i  20  net. 
The  same.     Elementary  Course.     423  pp.     12010.     $1.00  net. 

7.  General  Biology.     By  Prof.   W.  T.  SEDGWICK,  of  Massachusetts  Institute  of 

Technology,  and  Prof.  E.  3.  WILSON,  of  Columbia  College.     (Revised  and 
enlarged^  1896.)     231  pp.     8vo.     $1.75  net. 

8.  Psychology.     By  WILLIAM  JAMES,  Professor  in  Harvard  College.    Advanced 

Course.     689  +  7°4  pp.     8vo.     2  vols.     $4.80  net. 
The  same.     Briefer  Course.     478  pp.     12010.     $1.60  net. 

9.  Physics.     By  GEORGE  F.  BARKER,  Professor  in  the  University  of  Pennsylva- 

nia.    Advanced  Course.     902  pp.     8vo.     $3.50  net. 

10.  Geology.     By  THOMAS  C.  CHAMBERLIN  and  ROLLIN  D.  SALISBURY,  Professors 

in  the  University  of  Chicago.     (/«  Preparation.) 
1  1.     Finance.      By  HENRY  CARTER  ADAMS,  Professor  in  the  University  of  Michi- 

gan.    Advanced  Course,     xiii  -f-  573  pp.     8vo.     $3.50  net. 
in,  1900  (l) 


HENRY  HOLT  &    CO.'S  WORKS   ON  SCIENCE. 

Allen's    Laboratory  Exercises   in    Elementary  Physics.       By  CHAS.   R. 

ALLEN,  of    the    New    Bedford,    Mass.,    High    School.     Pupils' 

Edition,  x -)- 209  pp.,   8oc.,  net.      Teachers'  Edition,   $1.00,  net. 
Arthur,  Barnes,  and  Coulter's  Handbook  of  Plant  Dissection.     By  Prof. 

J.  C.  ARTHUR,  of  Purdue  Univ.,   Prof.  C.  R.  BARNES,  of   Univ. 

of   Chicago,    and    Pres.   JOHN   M.    COULTER,   of    Lake    Forest 

Univ.     xi  -(-  256  pp.     $1.20,  net. 
Atkinson's    Elementary    Botany.     By    Prof.    GEO.    F.    ATKINSON,   of 

Cornell.      Fully  illustrated,      xxiii  -f-  441  pp.      $1.25,  net. 

Atkinson's  Lessons  in  Botany.     Illustrated.     365  pp.     $1.12,  net. 
Barker's  Physics.     See  American  Science  Series. 

Barnes's  Plant  Life.  By  Prof.  C.  R.  BARNES,  of  University  of 
Chicago.  llhistrat,d.  x -f- 428  pp.  $1.12,  «/?/. 

Barnes's  Outlines  of  Plant  Life.     Illustrated.     308  pp.     $1.00,  net. 

Beal's  Grasses  of  North  America.  For  Farmers  and  Students,  By 
Prof.  W.  J.  BEAL,  of  Mich.  Agricultural  College.  Copiously 
Ill'd.  8vo.  Vol.  I.,  457  pp.  $2.50,  net.  Vol.  II.,  707  pp.  $5,  net. 

Bessey's  Botanies.     See  American  Science  Series. 

Black  and  Carter's  Natural  History  Lessons.  By  GEO.  A.  BLACK,  and 
KATHLEEN  CARTER.  (For  the  very  young.)  98  pp.  soc.,  net. 

Britton's  Manual  of  the  Flora  of  the  Northern  States  and  Canada.     By 

Prof.  N.  L.  BRITTON,  Director  of  N.  Y.  Botanical  Garden. 

Bumpus's  Laboratory  Course  in  Invertebrate  Zoology.  By  H.  C.  BUMPUS, 
Professor  in  Brown  University.  Revised.  157  pp.  $i,  net. 

Cairns's  Quantitative  Chemical  Analysis.  By  FRED'K  A.  CAIRNS.  Re- 
vised and  edited  by  Dr.  E.  WALLER.  417  pp.  8vo.  $2,  net. 

Champlin's  Young  Folks'  Astronomy.  By  JOHN  D.  CHAMPLIN,  Jr., 
Editor  of  Champlin's  Young  Folks'  Cyclopaedias.  Illustrated, 
vi  -f-  236  pp.  i6mo.  48c.,  net. 

Congdon's  Qualitative  Analysis.  By  ERNEST  A.  CONGDON,  Professor 
in  Dtexel  Institute.  64  pp.  Interleaved.  8vo.  6oc.,  net. 

Crozier's  Dictionary  of  Botanical  Terms.     202  pp.     8vo.     $2.40,  net. 

Hackel's  The  True  Grasses.  Translated  from  "Die  natlirlichen 
Pflanzenfamilien "  by  F.  LAMSON-SCRIBNER  and  EFFIE  A. 

SOUTHWORTH.       V  -[  •  228  pp.       8vo.       $1.50. 

Hall's  First  Lessons  in  Experimental  Physics.  For  young  beginners, 
with  quantitative  work  for  pupils  and  lecture-table  experiments 
for  teachers.  By  EDWIN  H.  HALL,  Assistant  Professor  in  Har- 
vard College,  viii  4-  120  pp.  I2mo.  6sc.,  net. 

Hall  and  Bergen's  Text-book  of  Physics.  By  EDWIN  H.  HALL,  Assist- 
ant Professor  of  Physics  in  Harvard  College,  and  JOSEPH  Y. 
BERGEN,  Jr.,  Junior  Master  in  the  English  High  School,  Bos- 
ton. Greatly  enlarged  edition.  596  pp.  I2mo.  $1.25,  net. 

Postage  8%  additional  on  nst  books.     Descriptive  list  free. 
in,  1900 


HENRY  HOLT  &    CO.'S   WORK'S   ON  SCIENCE. 

Hertwig's  General" Principles  of  Zoology.  From  the  Third  Edition  of 
Dr.  Richard  Hertwig's  Lehrbuch  der  Zoologie.  Translated  and 
edited  by  GEORGE  WILTON  FIELD,  Professor  in  Brown  Univer- 
sity. 22*6  pp.  8vo.  $1.60  net. 

Howell's  Dissection  of  the  Dog.  As  a  Basis  for  the  Study  of  Physi- 
ology. By  W.  H.  HOWELL,  Professor  in  the  Johns  Hopkins 
University.  100  pp.  8vo.  $1.00  net. 

Jackman's  Nature  Study  for  the  Common  Schools.  (Arranged  by  the 
Months.)  By  WILBUR  JACKMAN,  of  the  Cook  County  Normal 
School,  Chicago  111.  448  pp.  $1.20  net. 

Kerner  &  Oliver's  Natural  History  of  Plants.  Translated  by  Prof.  F. 
W.  OLIVER,  of  University  College,  London.  410.  4  parts. 
With  over  1000  illustrations  and  16  colored  plates.  $15.00  net. 

Kingsley's  Vertebrate  Zoology.  By  Prof.  J.  S.  KINGSLEY,  of  Tufts 
College.  Illustrated.  439  pp'  8vo.  $3.00  net. 

Kingsley's  Elements  of  Comparative  Zoology.    357pp.    i2mo.    $i.20»rf. 

Macalister's  Zoology  of  the  Invertebrate  and  Vertebrate  Animals.  By 
ALEX.  MACALISTER.  Revised  by  A.  S.  PACKARD.  277  pp. 
i6mo.  8oc.  net. 

MacDougal's  Experimental  Plant  Physiology.  On  the  Basis  of  Gels' 
Pflanzenphysiologische  Versuche.  By  D.  T.  MAcDoUGAL,  Uni- 
versity of  Minnesota,  vi  -f-  88  pp.  8vo.  $1.00  net. 

Macloskie'S  Elementary  Botany.  With  Students'  Guide  to  the  Exam- 
ination and  Description  of  Plants.  By  GEORGE  MACLOSKIE, 
D.Sc.,  LL.D.  373  PP-  $i-3Q  net. 

McMurrich's  Text-book  of  Invertebrate  Morphology.  By  J.  PLAY  FAIR 
McMuRRiCH,  M.A.,  Ph.D.,  Professor  in  the  University  of  Cin- 
cinnati, vii  -f-  661  pp.  8vo.  New  Edition.  $3.00  net. 

McNab's  Botany.  Outlines  of  Morphology,  Physiology,  and  Classi- 
fication of  Plants.  By  WILLIAM  RAMSAY  McNAB.  Revised  by 
Prof.  C.  E.  BESSEY.  400  pp.  i6mo.  8oc.  net. 

Martin's  The  Human  Body.     See  American  Science  Series. 

*Merriam's  Mammals  of  the  Adirondack  Region,  Northeastern  New 
York.  With  an  Introduction  treating  of  the  Location  and 
Boundaries  of  the  Region,  its  Geological  History,  etc.  By  Dr. 
C.  HART  MERRIAM.  316  pp.  8vo.  $3.50  net. 

Newcomb  &  Holden's  Astronomies.     See  American  Science  Series. 

Nicholson  &  Avery's  Exercises  in  Chemistry.  By  Prof.  H.  H.  NICHOL- 
SON, University  of  Nebraska,  and  Prof.  SAMUEL  AVERY,  Uni- 
versity of  Idaho.  134  pp.  6oc.  net. 

*Noel's  Buz  ;  or,  The  Life  and  Adventures  of  a  Honey  Bee.  By 
MAURICE  NOEL.  134  pp.  $1.00. 

Noyes's  Elements  of  Qualitative  Analysis.  By  WM.  A.  NOYES,  Pro- 
fessor in  the  Rose  Polytechnic  Institute.  91  pp.  8vo.  8oc.  net. 

Packard's  Entomology  for  Beginners.  For  the  use  of  Young  Folks, 
Fruit-growers,  Farmers,  and  Gardeners.  By  A.  S.  PACKARD. 
xvi -1-367  pp.  Third  Edition,  Revised.  $1.40  net. 

«H,    IQOO 


HENRY  HOLT  &    CO.'S    WORKS   ON  SCIENCE. 


Packard's  Guide  to  the  Study  of  Insects,  and  a  Treatise  on  those 
Injurious  and  Beneficial  to  Crops.  For  Colleges,  Farm-schools, 
and  Agriculturists.  By  A.  S.  PACKARD.  With  15  plates  and 
670  wood-cuts.  Ninth  Edition.  715  pp.  8vo.  $4.50  net. 

Outlines  of  Comparative  Embryology.  Illustrated.  243pp.  8vo. 

$2.00  net. 

Zoologies.     See  American  Science  Series. 

Peabody's  Laboratory  Exercises  in  Anatomy  and  Physiology.  By  JAS. 
EDWARD  PEABODY,  of  the  High  School  for  Boys  and  Girls, 
New  York.  x-|-79pp.  Interleaved.  I2mo.  6oc.  net. 

Perkins's  Outlines  of  Electricity  and  Magnetism.  By  Prof.  CHAS.  A 
PERKINS,  of  the  University  of  Tennessee.  277  pp.  i2mo. 
$1.10  net. 

Pierce's  Problems  in  Elementary  Physics.  Chiefly  numerical.  By  E. 
DANA  PIERCE,  of  the  Hotchkiss  School.  194  pp.  6oc.  net. 

*  Price's  The  Fern  Collector's  Handbook  and  Herbarium.    By  Miss  SADIE 

F.  PRICE.     72  plates,  mostly  life-size,  with  guide.     410.     $2.25. 
Randolph's  Laboratory  Directions  in  General  Biology.    163  pp.    Soe.  net. 
Remsen's  Chemistries.     See  American  Science  Series. 
Scudder's  Butterflies.    By  S.  H.  SCUDDER.    322pp.    i2mo.   $i.2o»;/. 
Brief  Guide  to  the  Commoner  Butterflies,     xi  -f  206  pp.     $1.25. 

The  same.     With  21  plates,  containing  qj  illustrations.    $1.50. 

The   Life  of  a  Butterfly.     A  Chapter  in  Natural  History  for  the 

General  Reader.     By  S.  H.  SCUDDER.     186  pp.     i6mo.     8oc.  net. 
Sedgwick  and  Wilson's  Biology.     See  American  Science  Series. 
*Step's  Plant  Life.    Popular  Papers.    Illustrated.    218  pp.    $100  net. 
Torrey's  Elementary  Studies  in  Chemistry.    By  JOSEPH   TORREY,    Jr., 

Instructor  in  Harvard.     487  pp.     $1.25  nef- 
Underwood's  Our  Native  Ferns  and  their  Allies.     By  Prof.  LUCIEN  M. 

UNDERWOOD,  of  Columbia.     156  pp.     $1.00  net. 
Underwood's  Moulds,  Mildews,  and  Mushrooms.  ///'</.   236pp.  $i.$onef. 
Williams's   Elements  of    Crystallography.     By   GEORGE    HUNTINGTON 

WILLIAMS,  late   Professor    in    the    Johns    Hopkins   University. 

x  -}-  270  pp.     Revised  and  Enlarged.     $1.25  net. 
Williams's   Geological    Biology.    An     Introduction  to  the  Geological 

History  of  Organisms.     By  HENRY  S.  WILLIAMS,  Professor  of 

Geology  in  Yale  College.     8vo.     395  pp.     $2.80  «*•/.  * 
Woodhull's  First  Course  in  Science.     By  JOHN  F.  WOODHULL,  Pro- 
fessor in  the  Teachers'  College,  New  York  City. 

/.     Book  of  Experiments,    xiv  -\-  79  pp.    8vo.    Paper.    5oc.  net. 
II.      Text-Book,     xv  +  133  Pp.     I2mo.      Cloth.     6^c.  net. 
Woodhull  and  Van   Arsdale's    Chemical  Experiments.     An    elementary 

manual,    largely  devoted    to  the    chemistry  of    every-day    life. 

Interleaved.      136  pp.      6oc.  net. 
Zimmermann's  Botanical  Microtechnique.     Translated  by  JAMES  ELLIS 

HUMPHREY,  S.  C.     xii -f  296  pp.     8vo.     $2.50  net. 

HENRY  HOLT  &  CO.,  29  WEST  230  ST.,  NEW  YORK. 
in, 1900 


THE  UNIVERSITY  OF  CALIFORNIA  LIBRARY 

I  T 


iiiii 


